hello everybody my name is Iman welcome back to my YouTube channel today we're going to start chapter 2 for General chemistry 1 this chapter is titled atoms molecules and ions now in this chapter we're going to cover the following objectives first we're going to start with a discussion on fundamental chemical laws here we're going to define the law of conservation of mass the law of definite proportions and the law of multiple proportions next we're going to learn about Dalton's atomic theory so in 1808 Dalton published a new system of chemical philosophy in which he presented his theory of atoms that we're going to go over then our third objective is a continuation of some of the earlier ideas of the basic structure of an atom we're going to compare that to our current understanding of an atom before we move into objective four which will further build on on the modern view of atomic structure here we're going to discuss what is an atom made of and how do the atoms of the various elements differ in addition we're going to cover a lot of important terminology here in objective 4 we're going to talk about atomic number mass number atomic weight atomic mass unit moles and molar mass then the fifth objective is on molecules and ions and here we're also going to learn the difference between coent and ionic bonding then we're going to move to the sixth objective which is an introduction on the periodic table we will also cover periodic table Trends this is going to require us to understand atomic radius effective charge ionization energy electr negativity as well as electron affinity and then last but not least we get to move on to our last and final objective we're going to learn how to name simple compounds now with that let's go ahead and get started but before we even dive into objective one I really want to set the stage for this chapter I really like to do this because having a little bit of a background story really helps to hopefully peque your interest into the topics we're going to cover it and why we want to cover them chemistry has been important since ancient times the Greeks were the first to try to explain why chemical changes occur and by about 400 BC they had proposed that all matter was composed of four fundamental substances fire earth water and air only the Avatar mastered all four elements but when the world needed him most he vanished and so did this idea but the Greeks they had considered a lot of questions that really lays the basis for some of the for for chemistry they asked whether matter is continuous and whether it was infinitely divisible into smaller pieces or not they had no experiments that supported their claims but they had these questions they had these thoughts and ideas and actually the next 2,000 years of chemical history were dominated by a pseudo science called Alchemy some Alchemists were fakes who were just really obsessed with the idea of just turning cheap metal into gold but many Alchemists were indeed serious scientists and this period did see some important advances The Alchemist discovered several elements and they learned to prepare the mineral acids Etc now the first quote unquote chemist to perform truly quantitative experiments was Robert Bole who measured the relationship between the pressure and volume of air and we're going to talk about Robert Bole a lot more when we get to our chapter on gases now all his ideas weren't correct but he was an excellent science scientist and he did lay some groundwork that was extremely important and after him a lot of interest in the 17th and 18 18th century actually centered around the phenomena of combustion the phenomena of combustion really evoked intense interest in the 17th and 18th century and by the late 18th century combustion had been studied extensively and the gases like carbon dioxide nitrogen hydrogen and oxy had been discovered and it was Antoine lavier forgive my pronunciation I'm not French and I suck at pronouncing French names but he was a French chemist who finally explained the true nature of combustion and his precise measurements in chemical reactions demonstrated a really important concept that's going to allow us to get started here with our first objective on fundamental chemical laws and that was the law of conservation of mass he showed that combustion and respiration involved oxygen a gas he named um conservation of mass is really important to understand and it essentially states that in a closed system the total mass of the reactants which are the substances that start a chemical reaction equals the total mass of the product products are the substances that are formed by the reaction and so in simple terms mass is neither created nor destroyed in chemical reactions it's simply transformed we can use this example right here to kind of motivate this idea right what we see here is methane CH4 it reacts with oxygen gas O2 to form carbon dioxide CO2 and water H2O all right this is going to be used as an example to illustrate the law of conservation of mass so we have one molecule of methane CH4 that's going to be 16 G this is just adding up the individual weights of all the atoms that make up CH4 here carbon 12 hydrogen 1 and there's four of them so total of 16 G then we have two2 molecules all right O2 that's two oxygens each oxygen is 16 for a total of 32 that's for one2 molecule and if you have two that's 64 4 G so for the reactants we have a total of 80 g the total mass before the reaction is 80 g that's 16 G of methane plus 64 G of oxygen after the reaction is complete these reactants are going to form carbon dioxide and two water molecules now if you add up the weights of those you're going to get 44 G for 1 CO2 and 36 G for 2 H2O molecules and lo and behold 44 + 36 = 80 g the mass of these products combined will also be 80 g demonstrating the law of conservation of mass now this law holds for All chemical reactions showing that mass is conserved throughout the process all right no Mass lost or gained in the reaction now building on Antoine's principles Joseph Prost another French chemist estab established the law of definite proportions how let's talk about that next he established the law of definite proportion by demonstrating that compounds always contain the same proportion of elements by mass his findings actually influenced John Dalton who proposed that Adams as indivisible particles make up elements this atomic theory explained why compounds always have consistent relative masses of their constituent elements but Dalton discovered another principle that convinced him even more of the existence of atoms and that was the law of multiple proportions the law of multiple proportions this said when two elements form a series of compounds the ratios of the masses of the second element that combined with 1 G of the first element can always be reduced to small whole numbers so let's kind of demonstrate this with a problem here we have data that was collected for several compounds of nitrogen and oxygen so you have compound a b and c and here you have the mass of nitrogen that combines with one gam of oxygen so in compound a you have 1.75 G of nitrogen that combines with 1 gram of oxygen in compound B you have 0.875 gram of nitrogen that com that combines with one gram of oxygen and C you have 0.4375 G of nitrogen that combines with 1 G of oxygen now to demonstrate the law of multiple proportions we can start comparing these compounds we can compare compound a to B for example the mass of nitrogen combining with 1 G of oxygen in compound a is 1.75 G and that for B is 0.8750 G all right this is exactly 2 to one all right this gives us a 2 over one relationship and so the mass of nitrogen combining with 1 gam of oxygen in compound a is exactly double that in compound B now we could do the same we can compare B and C for example this is going to be 0.8750 G over 0.4375 G and this also gives us a 2 over one relationship so it is also double here the mass of nitrogen combining with 1 G of oxygen in compound B is exactly double that in compound C we can do another comparison we can do a and C this is going to be 1.750 G over 0.4375 G this is going to give us a 4 over 1 relationship so the mass of nitrogen combining with 1 G of oxygen in compound a is exactly four times that in compound C and what you notice is these ratios between the different compounds 2: 1 2: 1 4:1 are all small whole numbers which which perfectly illustrates the law of multiple proportions and just to kind of summarize that I want to restate the definition once more the law of multiple proportions is a fundamental principle in chemistry stating that when two elements combined to form more than one compound the masses of one element that combined with a fixed mass of the other are in a ratio of small whole numbers now with that discussion we can go ahead and move into objective two which is on Dalton's atomic theory in 1808 Dalton published a system of chemical philosophy in which he presented his theory of atoms and this is what he wrote first he said each element is made up of tiny particles called atoms second the atoms of a given element are identical and the atoms of different elements are different in in some fundamental way or ways three chemical compounds are formed when atoms of different elements combine with each other a given compound always has the same relative number and types of atoms and last but not least chemical reactions involve reorganization of the atoms so changes in the way they are bound together the atoms themselves are not changed in a chemical reaction John Dalton really viewed atoms as the indivisible building blocks of matter he viewed them kind of like billiard balls atoms of different elements simply differed from each other by mass and nothing else they could not be subdivided any further it clearly did not account for the subatomic particles that make up atoms that was discovered later and this is the perfect segue into objective three where we want to talk about all the early ideas and experiments that aim to understand and characterize atoms the concept of atoms is really interesting and inevitably scientists began to wonder about the nature of the atom what is an atom made of how do the atoms of the various elements differ so let's start that conversation we now know that an atom is the smallest unit of ordinary matter that forms a chemical element it's composed of a nucleus that's made out of protons positively charged and neutrons which have no charge electrons which are negatively charged orbit the nucleus the protons and neutrons Define the mass of the atom while the arrangement and movement of electrons determines the atoms chemical properties and reactivity how did we reach this conclusion right it was the cumulative work of many scientists that has allowed us to understand the basic structure of an atom we just made mention of John Dalton who proposed his atomic theory and how he viewed atoms as sort of like billiard balls atoms of different elements simply differed from each other by mass and nothing else and they could not be subdivided any further so this is kind of our starting point here it wasn't until JJ Thompson's work with caid Ray tubes in 1898 that negatively charged particles called electrons were found to be present in atoms and so Thompson's model is usually called The Plum putting model you have these scattered negative charges around um in an atom then in 1911 Ernest Rutherford further refined our view of the atom by showing that although atoms consist of mostly empty space at the center of every atom is an extremely small and dense nucleus which is positively charged he was able to make these conclusions based off of his gold foil experiment so then we had Rutherford model Center nucleus positively charged and then the negative charges the electrons kind of surrounded this extremely small and dense nucleus then bour came along and he developed his model which suggests that negative charge orbits the nucleus of an atom which is made out of protons and neutrons there is more to that truly as we can't really pinpoint where electrons are we can only really talk about electron density probability where the electrons most likely to be found but we're going to dive into that later so essentially we know that atoms have a nucleus made out of protons and neutrons and surrounding that are electrons in our quantum mechanical model our latest model the electrons kind of exist in a cloud of probability around the nucleus we don't really want to think about electrons as quote unquote orbiting the nucleus but more so as existing as a cloud of probability electron density probability around the nucleus we're not going to get too much into that until maybe later cuz it is a complicated topic but now that we understand that I do want to focus on the subatomic particles in more detail we've said that an atom is the smallest identifiable unit of an element so in our fourth objective we're going to focus on that modern view of atomic structure we have an atom we talked about how it contains a nucleus with protons these are subatomic particles with a positive charge neutrons subatomic particles with no electric charge they make up the nucleus and then we have electrons subatomic particles with negative charge now in addition to that we should know some quantifiable information about protons neutrons and electrons here we have a table with some of that important information proton the mass in kilograms all right 1. 672 * 10us 27 kilograms or the mass in AMU atomic mass units we'll talk about that in a second is about 1727 the charge for a proton is 1.62 * 10us 19 and the relative charge is + one so this is the charge in kums neutrons here is the mass in kilogram the mass in AMU it has zero charge and then we have the electron it's the smallest of the three as you can obviously tell by the mass 9109 * 10- 31 kg in amus it's 0.548 the charge is the same as the proton but with a negative sign so protons had positive charge electrons have a negative charge now something to know about electrons is that they move around the nucleus at varying distances away from the nucleus we're going to get into a lot of the details about this a little later on but some few points that I want you to keep in mind let them marinate in your subconscious until we get to it later on but the electrons that are closer to the nucleus are at lower energy levels and those that are further out in higher shells are going to have higher energy and while we're going to revisit this in more detail we should know that the electrons furthest away away from the nucleus are going to have the strongest interactions with the surrounding environment you can think about the electrons further away from the nucleus as being kids that are not as close to their parents and therefore they have a stronger interaction with their surroundings their friends their environment their school Etc just as a analogy if you will they're going to have the strongest interactions with the surrounding environment and the weakest interactions with the nucleus since it's further away these electrons are called veence electrons they are the electrons that are going to participate in chemical reactions and they are more likely to participate in bonds and this is a central idea that you're going to carry even into organic chemistry now that we understand the atom let's talk about a couple of important terms related to the atom so atomic structure is defined by three critical numerical descript scriptors we have atomic number mass number and atomic weight so typically if you're looking at an element on the periodic table for example you'll see something like this you have the elemental symbol for oxygen for example the name of the element at the top here is your atomic number at the bottom here is your mass number but in addition you might see this information communicated in other ways for example in this format where you have your Elemental Sy and then at the top left corner you have your mass number and your bottom left corner your atomic number so oxygen could look like this on the periodic table but sometimes that same information can be communicated like this in writing regardless of the format I want to Define some of these terms so atomic number denoted by the symbol Z is the count of protons in the nucleus of an atom this number is fundamental because it determines the chemical identity of the atom each element on the periodic table is characterized by a unique atomic number then you have the mass number so let's do this in a different color the mass number denoted by a so you could see that here highlighted in red this is the total number of protons and neutrons in an atom's nucleus neutrons and protons are collectively known as nucleons and their sum gives us the mass number now the difference between the mass number and the atomic number is the number of neutrons in the nucleus it's important to note that the mass number is not the same as the atomic weight atomic weight is a more nuanced concept it is the weighted average of the masses of all naturally occurring isotopes of an element and it's measured in atomic mass units since isotopes of an element have same number of protons but varying numbers of neutrons they also have different Mass numbers but of course they'll have the same atomic number so atomic weight takes into account the relative abundance of each isotope in nature in the calculation and that kind of makes it important for us to really Define what an isotope is so let's quick Define that and then let's go back to this topic of atomic weight to make sure that we are all on the same page Isotopes are variant of elements that have the same number of protons but different number of neutrons and that leads to different Mass numbers they're significant in both natural processes and technological applications so for instance stable isotopes are used in medical Diagnostics whereas radioactive isotopes can be applied in treatment or to help date archaeological finds or as tracers in biochemical research understanding atomic structure and Isotopes is really important for grasping the bigger picture of chemical behavior so with that being said let's go ahead and go back to atomic weight now that we have defined isotopes all right we defined atomic weight and we have it right here for us for our convenience as the weighted average of the masses of all naturally occurring isotopes of an element measured in atomic mass units so let's really break this down because this calculation is not super straightforward it takes into account all right it has to take into account the fact that most elements exist naturally as a mixture of Isotopes each with its own mass number and Natural Abundance and so to calculate atomic weight scientists consider both the mass and the Natural Abundance of each isotope and here how here's how that works you have to consider one all right we're going to make a little list over here one you have to consider the isotope masses all right first the exact mass of each is of the element is determined these masses are close to whole numbers but are slightly different all right then you have to consider two Natural Abundance Natural Abundance you have to know each Isotopes Natural Abundance this is the percentage of each isotope that occurs naturally on Earth and then you can go ahead and calculate that weighted average the atomic weight is then calculated at by multiplying the mass of each isotope by its natural uh abundance and then adding these values together the result is a weighted average that represents the average mass of all isotopes that occur naturally now the best way to understand this is by doing a quick example so what I'm going to do is I'm going to scroll down here really quickly and I have a problem that relates to what we just discussed about atomic weight that we're going to do together so in this problem we say we have an element this element has two isotopes isotope a has a mass of 10 AMU and we're going to cover what AMU is in a second but let's just say that this is a measure of mass here all right isotope a with a mass of 10 AMU and a Natural Abundance of 90% And then we have isotope B with a mass of 11 AMU and a Natural Abundance of 10% what is the atomic weight So based off of our verbal description of atomic weight we said that it would be calculated like so we have our first isotope isotope a okay it's 10 AMU in mass and it has a Natural Abundance of 90% And we want to express 90% as a decimal so that's 0.90 okay that's for isotope a and we're going to have to add this to isotope B all right isotope B 11 AMU and it has a Natural Abundance of 10% so 0.10 all right the atomic weight can then be calculated as follows we just follow through with this mathematical expression right here plug this into a calculator follow pem rules okay and what we're going to get is that this is going to equal to 9 plus oh let me use my pen 9 + 1.1 and that gives us 10.1 AMU and so the atomic weight of this element all right that has two isotopes is 10.1 AMU all right so again the key to understanding atomic weight is to recognize that it's an average that reflects isotopic composition and abundance it's not the mass of a single atom but an average mass of all atoms of that element as they are found in nature wonderful now let's scroll back up because we skipped a few things to get to that problem that we still want to go ahead and cover so we've talked about atomic number mass number and atomic weight there are a couple more Concepts that are going to be super crucial for us to understand before we move on and those are the concepts of mass atomic mass unit and moles so let's go ahead and cover those now let's start with atomic mass unit the atomic mass unit is a standard unit of mass that quantifies the mass of atoms or molecules and it's defined as 12th the mass of a carbon 12 atom so let me repeat that it's defined as 1 12th the mass of a carbon 12 atom this is an isotope of carbon with six protons and six neutrons it's stable and its abundance makes it a standard for measuring atomic masses one atomic mass unit is approximately 1.66 * 10- 27 - 27 I want to make that clear kilog now the atomic mass of an element that's usually found on the periodic table is expressed in atomic mass unit and it represents the average mass of all the Isotopes of that element taking into account their relative abundance the next thing I want to talk about is moles a mole is a unit that measures the amount of substance so like one dozen signals that you have 12 things right if you say you have one dozen eggs you have 12 eggs you have one dozen books you have 12 books one dozen magazines you have 12 magazines same thing for a mole it is a unit that measures the amount of a substance one mole of any substance contains exactly 6.022 * 10 23 entities this could be atoms molecules ions Etc this number is is known as avagadro's number the mole allows chemist to count particles by weighing them and the mass of one mole of a substance is equal to its molecular or atomic mass in grams and last but not least I want to talk about molar mass this is the mass of one mole of a substance it's expressed in units of gam per mole now that we've Define these three terms atomic mass unit moles molar mass how do we connect them all together and by that I mean how do we convert from one form into another this is going to be really important in general chemistry so let's talk about it let's talk about how we can connect these Concepts let's say to start off you know the number of particles you know the number of particles and you want to convert that to moles how do you go from number of particles to moles you do this by dividing by avagadro's number all right you do this by dividing by avagadro's number so let me demonstrate this a little bit better all right you divide divide by avagadro's number okay that's going from number of particles to moles and that's kind of written up here at the top as well okay what about if you know the number of moles but now you want number of particles right so you're going in the opposite direction now you know moles you want number of particles here now you're going to have to multiply by avagadro's number fantastic so we got that conversion number of particles to moles or moles to number of particles we know that the secret lies in avagadro's number phenomenal what if you know moles all right and you want to convert to mass you want to go from moles to mass how do you do that in order to do that you're going to multiply by molar mass so you're going to multiply by G per mole and if you want to go the other way if you know the mass and you want the moles you're going to divide by molar mass so you're going to think mole over grams good so far so good here's something else what if you know the mass and you want to convert all right you want to go from Mass to number of particles what do you do now all right this is going to involve two steps and it's the two steps that we've talked about going in the same direction first you're going to divide by molar mass all right so Step One is divide by molar mass and then the second step is to multiply by avagadro's number multiply I can't spell by avagadro's number so if you're going from Mass to number of particles two steps you divide by molar mass then you multiply by avagadro's number okay last thing that we want to cover now is you guessed it what if we want to go from number of particles to mass again two steps First Step you divide by avagadro's number and then the second step you multiply by molar mass so this is kind of like a flowchart a road map of how to convert between number of particles moles and mass let's do a quick example to make sure that we understand this concept fully so this is the problem right here suppose you have 3.11 * 10 23 particles of argon what is the mass of the Argon sample so here's our goal we know the number of particles we want to go to mass so this is the path we want to follow this one right here this is going to require us to divide by avagadro's number and then multiply by molar mass now the molar mass of argon is going to be 39.95 G per mole we're just going to go ahead and kind of set up this dimensional analysis we're starting off what do we know we know that we have let's pick a color let's do green 3.11 * 10^ 23 particles of argon all right we want to convert this to mass ultimately first step is divide by avagadro's number so avagadro's number is 6.022 * 10 to the 23 particles all right is equal to 1 mole phenomenal so this is our first conversion okay notice how mole by avag godra number that's exactly what we have here mole by avagadro's number we have particles here because that's what we're dealing with particles of argon and the units here cancel out good now what's the next step the next step is to multiply by molar mass so what's the molar mass of argon is just 39.95 G per mole in notice again mole cancels out and that means our final answer is going to have units of grams and that's exactly what we want cuz we're trying to figure out what the mass of this argon sample is and then you just go ahead and calculate this you should get about 19.98 G of argon and there we have it we did it now in summary we've talked about a lot here in objective 4 I want to scroll back up and just show you how much we've talked about we talked about protons neutrons and electrons the subatomic particles that make up an atom we talked about atomic number and mass number and how to identify it in different forms in different writing styles if you will and then we had a little bit of a tangent on how mass number is not equal to atomic weight and so we were motivated to talk about atomic weight which also required us to talk about Isotopes and we learned how to calculate atomic weight and that atomic weight is defined as the weight weighted average of the isotope masses of the elements that are naturally occurring Isotopes and then we took that and we contined to talk about a couple of other important terminology atomic mass unit moles and molar mass and we learned how to convert between different units going from number of particles to moles to mass and back and forth and in between and in the process we did two example problems as well now we can confidently move into our fifth objective our fifth objective is about molecules and ions something important to remember is that in ordinary chemical reactions the nucleus doesn't change but the number of electrons that an atom possesses can change the only time that the number of protons or neutrons possessed by an atom change is during nuclear reactions we're not talking about that here now it's safe to say that the movement of electrons constitute reactions and this will be the focus of most discussions with reactions in general and organic chemistry now it's important to pay close attention to whether a particle is charged or not if a particle is a charged ion or a neutral atom because an ion possesses properties that are quite different from those of a neutral molecule an ion is a charged particle resulting from the fact that the number of protons is not equal to the number of electrons so when the number of protons equals the number of electrons we're dealing with a neutral molecule but when the number of particles does not equal the number of electrons then we are dealing with ions there are two types of ions we have c i and anion a positively charged atom is called a cat c i and a negatively charged atom is called an anion so for cat I they're positively charged because the number of protons is greater than the number of electrons so you've lost an electron so hence catons are positively charged anions on the other hand you have a greater number of electrons than protons you've gained an electron so they are negatively charged again it's important to pay attention to whether a particle or molecule is charged or not because an ion possesses properties that are quite different from those of a neutral atom understanding ions allows us to discuss molecules and ionic compounds molecules are distinct electrically neutral groups of atoms that are held together by chemical bonds specifically Co Co valent bonds where atoms share electrons we'll get into the details of that here just in the next page the structure and properties of a molecule they're determined by the specific atoms it contains and how those atoms are bonded a molecule's formula indicates the types and numbers of atoms present using element symbols and subscripts to denote their ratios on the other hand in contrast ionic compounds they consist mainly of oppositely charged ions that are held together by ionic bonds these are electrostatic forces between the ions again we're also going to get into that here in the next page these compounds typically form highly ordered three-dimensional Crystal lce structures the formula of an ionic compound represents the smallest neutral unit of a compound indicating the ratio of catons to anion so again the ratio of catons positively charged ions to anion negatively charged ions so for example a good a good example is table salt table salt is sodium chloride na this is an ionic compound comprising of a one: one ratio of sodium cat ions two chloride annion and they don't exist like this like a typical traditional calent bond this is an ionic compound and so it exists in a three-dimensional crystal lattice where you have these alternating bonds between sodium cations and chloride annion to form a specific crystal lce shape but that is an inorganic topic that we're not going to get into here now molecules molecules form when atoms are held together by Cove valent bonds this involves the sharing of electron pairs between atoms on the other hand ionic compounds form ionic bonds which occur when electrons are transferred from one atom to another resulting in positively charged ions catons and negatively charged ions anion that attract each other due to their opposite charges how can we distinguish between the two so now this is the perfect segue into talking about the different kinds of bonding calent and ionic bonding in order to do this we're going to have to define a really important topic called electr negativity electro negativity is a chemical property that describes the ability of an atom to attract electrons toward itself when it is part of a compound it's a dimension mless quantity often represented on the periodic table with values that typically range anywhere between 0.7 to 4.0 um which is a scale of lower to higher electronegative elements respectively so a lower number less electronegative higher number more electronegative now understanding electr negativity is vital because it helps us predict the type and the strength of bonds that are going to form between atoms now depending on your course here you might be required to memorize some values for electr negativities but in addition here in the next objective we're going to talk about periodic table and periodic table Trends and we're going to learn the trend for electro negativity so sometimes you can typically follow the trends to be able to follow the flowchart that we're going to discuss next to determine what kind of bonding is occurring between two atoms now the concept of bonding is really Central to chemistry it explains how atoms bind together to form molecules bonds can be broadly classified as ionic or as Cove valent and this depends on the electr negativity difference between the bonding atoms so let's go ahead and get started first with this Branch The Cove valent bonding Branch Cove valent bonding OCC URS when two atoms share a pair of electrons and that shared electrons allows each atom to attain the electron configuration of a noble gas leading to a more stable molecule coent bonds they can be polar or non-polar so let's talk about each of these nonpolar coent bonds these bonds form between atoms with similar Electro negativities usually the electro negativity difference is less than 0.5 since the electron pull is almost equal the electrons here in non-polar calent bonds are shared fairly equally so here's a good example if you have a carbon bonded to a carbon if you're trying to decide what kind of bond that is you would take the electr negativity value for carbon and the electro negativity value for carbon and you would subtract the two from each other now the electr negativity value for carbon is 2.6 essentially what you're doing here is 2.6 minus 2.6 that equals zero that is less than 0.5 so that is one way you can conclude that the bond between a carbon and a carbon atom is a non-polar coent Bond another one is whenever you have bonds between the same atoms it makes sense that their electro negativity difference is going to be zero and if it's zero then what you you have is a non-polar Co valent Bond let's do another example what about carbon and hydrogen carbon is 2.6 and hydrogen let's go back to our table hydrogen should have a it's not on here but I'll tell you hydrogen should have an electro negativity value of about 2.1 all right so about 2.1 so what we have is 2.6 minus 2.1 that's about 0.5 traditionally something on the border like this Ty LLY falls under non-polar calent bonding here anyways carbon hydrogen bonds those are nonpolar calent bonds then we have polar calent bonds okay these bonds occur when there's a moderate difference in electr negativities greater than 0.5 but less than 1.7 all right if the electro negativity differences Falls in this range then what you have is a polar coent bond the more Electro negative atom attracts the shared electrons more strongly what this leads to is a partial negative charge on that atom and a partial positive charge on the other one good example is carbon and oxygen carbon has an electro negativity value of 2.6 oxygen let's scroll back up here 3.5 so 3.5 minus 2.6 this is equal to 0.9 This falls in this range between 0.5 and 1.7 so the bond between carbon and oxygen is a polar coent bond now which one's more Electro negative oxygen is because its electro negativity value is 3.5 what does that mean that means oxygen pulls those electrons those shared electrons a little more towards it because it's more electronegative it likes those electrons just a little more than carbon what that results in is a partial negative charge on that oxygen since it's pulling those electrons towards it and a partial positive charge on the carbon so that is polar coal bonds okay so that is our Cove valent Branch what about ionic bonding ionic bonds when the difference in electr negativity is significant so think a electr negativity difference greater than or equal to 1.7 ionic bonding is likely in an ionic bond one atom usually a metal loses one or more electrons it becomes a positively charged ion while the other atom usually a non-metal gains those electrons and becomes a negatively charged ion and this electrostatic attraction between the oppositely charged ions forms a strong ionic bond and sodium chloride is a classic example of an ionic compound so if we look up here sodium has electr negativity value of one and chloride 3.2 so 3.2 minus1 that's a difference of 2.2 that is greater than 1.7 this is an ionic bond again sodium chloride great example sodium positive cation chloride anion so that is bonding for us we talked about electro negativity measure of the ability of an atom to attract electrons and we use the difference in electro negativity values between two atoms to determine what kind of bonding is occurring all right look at us learning So Much Chemistry let's move on to the sixth objective this objective is an introduction to the periodic table here we see a chemist's favorite tool the periodic table of elements we're going to continuously refer to the periodic table in this course and we're going to build our knowledge but to start let's talk a little bit of history in 1869 the Russian chemist Demitri Mel published the first version of the periodic table of elements and he showed that ordering the known elements according to atomic weight revealed a pattern of periodically recurring physical and chemical properties since then the periodic table has been revised using the work of physicist Henry Mosley to organize the elements based on increasing atomic number which is the number of protons in an element rather than atomic weight and using this revised table many properties of elements that had not yet been discovered could be predicted the periodic table it created a visual representation of the periodic law which states the chemical and physical properties of the elements are dependent in a periodic way upon their atomic numbers now the modern periodic table arranges the elements into periods or rows and into columns or groups now each period all right talking about periods each period is filled sequentially and each element in a given period has one more proton and one more electron than the element to its left so this atom right here has this element right here has one more proton and one more electron than the element before it and so on and so forth now as for groups groups contain elements that have the same electron configuration in their veence shell and they share similar chemical properties we're going to talk about electron configuration in a lot more detail in different chapter but with this preliminary information we want to talk a little bit about the different element types that we're going to see and discuss and there are three element types we're concerned about here metals non-metals and metalloids metals are found on the left side and in the middle of the periodic table they include the active metals the transition metals and also don't forget these two lonely rows at the bottom the lanides and the actinides metals are shiny solids except for mercury by the way which is a liquid under standard conditions they generally have high melting points and densities there are exceptions like lithium which has the density about half that of water metals have the ability to be deformed without breaking um the ability of metal to be hammered into shapes is called malleability and the ability to pull metals or Draw them into wires is called ductility many of the transition metals which are these Metals right here by the way are going to have two or more oxidation States so charges when forming bonds with other atoms and because the veence electrons of all metals are only loosely held to their atoms they're free to move which makes Metals really good conductors of heat and electricity so that's Metals For Us what about nonmetals non-metals are found predominantly on the upper right side of the periodic table so over here they have they're generally brittle in the solid state they show little or no metallic luster they are poor conductors of heat and electricity and all of these characteristics are manifestations of the inability of non-metals to easily give up electrons non-metals are also less Unified in their chemical and physical properties than metals are separating the metals and non-metals in this stair step group right here that I've attempted to draw in Black so the elements that surround this stair step right here these are called metalloids the metalloids are also called semi metals because they share characteristics with both metals and non-metals the electr negativities and ionization energies of the metalloids they're going to lie between metals and non-metals and their physical properties like density melting point boiling point they vary widely and can be a combination of metallic and non-metallic characteristics now what we're kind of equipped to do is begin our discussion on periodic trends but to have a productive conversation on periodic trends we're going to have to define a couple of really important words we're going to have to Define atomic radius effective charge ionization energy electr negativity and electron affinity first let's talk about atomic radius the atomic radius is a term that's used to describe the size of an atom it's defined as the average distance from the center of the nucleus to the outermost boundary of the electron cloud this average distance it's not a fixed physical boundary because remember electrons exist in a region of probability rather than in a specific orbital path however the definition I want you to keep in mind for now is that the atomic radius is the distance from the center of the nucleus to the outermost boundary of the electron cloud second we want to talk about effective nuclear charge effective nuclear charge is the actual amount of positive charge experienced by an electron in a multi-electron atom it is the charge that's felt by an electron when both the actual nuclear charge and the repulsive effects of other electrons are considered this is really important concept it's Quantified as the nuclear charge minus the shielding effect due to inner shell electrons now this can be estimated using Slater's rules we might talk about this a little more down the line but in short the effective nuclear charge is the net positive charge that veence electrons experience from the nucleus it is less than the actual number of protons in the nucleus because those inner electrons partially Shield the outer electrons from the full positive charge so in simple terms I suppose you can think of it as number of protons minus the core electrons that are getting in the way so that gives you kind of that actual amount of positive charge that's experienced by this electron in particular all the number of protons in here and all the electrons in the way then next up we have ionization energy this is a critical Atomic property that quantifies the energy required to remove an electron from a gaseous atom or ion this is an endothermic process meaning that energy is absorbed to overcome the electrostatic attraction between that negatively charged electron and the positively charged nucleus ionization energy is a key indicator of an element's reactivity and it is influenced by atomic radius as well as effective nuclear charge again it is the energy required to remove an electron next up is electr negativity we've briefly defined this earlier it is a measure of an atom's ability to attract and retain shared electrons in a chemical bond now electro negativity chemical property that describes the tendency of an atom to attract electrons towards itself it's influenced by element's atomic number as well as the distance that its veence veence elect Rons reside from the charged nucleus so again measure of the atom's ability to attract and bind with electrons in a chemical bond and then the last thing that we want to talk about is electron affinity this is the measure of the energy change when an electron is added to a neutral atom forming a negative ion so it's kind of the opposite of ionizing ation energy this is the energy required to remove an electron electron affinity is the amount of energy released when an electron is added to a neutral atom or molecule to form a negative ion now these concepts are extremely important in the study of chemical periodicity they each describe different but interrelated aspects of atomic behavior and now what we want to do is take these ideas and dive into periodic trends what is the trend for each of those properties going from left to right what about going from bottom to top let's start with atomic radius first what is the trend for atomic radius going from left to right on the periodic table the atomic radius is going to decrease this is due to the increase in the number of protons which enhances the effective nuclear charge thereby by drawing those veence electrons closer to the nucleus what about atomic radius going from bottom to top on the periodic table here it's going to decrease as you ascend a group fewer electron shells are present which reduces the shielding effect and it allows the nucleus to exert a stronger pull on the remaining electron shells next up effective nuclear charge what's the trend for Effective nuclear charge going from left to right it is going to increase as you move across a period the number of protons in the nucleus increases that should increase the nuclear charge however electrons are also being added to the same energy level so their ability to Shield each other from the nucleus that doesn't increase proportionately therefore the actual increase in positive charge felt by the electrons is going to be greater G thereby enhancing the effective nuclear charge so it increases going from left to right what about from bottom to top the effective nuclear charge is also going to increase going from bottom to top so moving up a group electrons in lower shells are going to Shield the outer electrons less effectively as you move up because the inner electron shells are going to be closer to the nucleus and thus EX exert a stronger pull on the outer electrons and so this results in an increase in the effective nuclear charge for the veence electrons since they experience a more significant attraction to the nucleus next up ionization energy what's the trend for ionization energy moving from left to right it is going to increase the electrons are held more tightly by the stronger effective nuclear charge and so it requires more energy for electrons to be removed and what about going from bottom to top for ionization energy it is going to increase moving upward fewer electron shells are present the electrons are going to be held closer to the nucleus and so increasing the energy is needed to ionize them then we have electronegativity what is the trend for electr negativity going from left to right it is going to increase due to the crater effective nuclear charge which allows atoms to attract bonding electrons more strongly and then for electr negativity going from bottom to top it is also going to increase because atoms higher up in the group have fewer electron shells resulting in a stronger attraction between nucleus and the veence electrons and then last but certainly not least we have electron affinity what is the trend going from left to right for electron affinity it is going to increase this means the atoms are more likely to accept additional electrons and this trend is because the atomic nucleus becomes more positively charged due to the addition of protons which increases the effective nuclear charge without a proportional increase in shielding effects since those additional electrons are being added to the same energy level and so as a result the added electrons feel a stronger attraction to the nucleus what about going from bottom to top electron affinity increases going from bottom to top so as the veence shell gets closer to the nucleus so with that we've covered our periodic table Trends and now we've actually made it to our last and final objective for this chapter naming in this section we're going to specify the most important rules for naming compounds other than organic compounds organic compounds is a worry for organic chemistry so we're not going to concern with ourselves with that here instead we're going to begin with the systems for naming inorganic binary compounds these are compounds that are composed of two elements and we're going to classify this into various types for easier recognition now the first thing to kind of set the stage is let's talk about how do we determine if a substance is ionic or molecular CU that's going to be really important for naming chemical compounds and knowing which rules to follow so if you're looking at a binary compound and both elements are non-metals then the compound is molecular water is a great example of this but if one add uh if one element is a metal and the other is a non-metal then the compound is ionic and here so sodium chloride is a really good example now in addition to that I also want you to make yourself accommodated please know these common monatomic cations and anions and their name so here we have a list of cations ions and their names H+ hydrogen L i+ that's lithium na+ that's sodium so on and so forth uh ag+ that's silver and then on the other side we have annion H minus hydride um CL minus chloride so on and so forth now something else that's important to keep in mind that'll make remembering some of these names easier is making flash cards and or recognizing that at least for the charges that are associated with each of these cat ions you can kind of extract that information from the periodic table so what I'm going to do is I'm going to scroll back up here where we had an image of the periodic table and what you noticed in the common names for monatomic cations specifically is any of the elements that you find here in group one are going to have plus one charges the metals that you're going to find here in group two are going to have plus two charges and usually some of these elements in group three um are going to have plus three charges like aluminum and gallium so keeping that in mind that at least makes it easier to associate the charge with the ion but of course also please know the names and the same goes for annion there's one more I want to add here just to keep in mind as we go through the lecture and we'll go over it again oh minus this is called hydroxide so also keep that in mind through this lecture because we're going to make mention of this annion now let's go ahead and get started with all that preliminary fluff we can get to talking first about binary compounds now type one binary ionic compounds they contain a positive ion a cat that's always written first in the formula and then it's followed by a negative ion or a annion in naming these compounds we're going to follow a couple of rules the first rule is that the cation is always named first and then the annion the annion is second now how do we name those different parts how do we name the cation how do we name the anion that's what the next rules about the second rule says a monatomic cat it takes its name from the name of the element so for example na+ that's just called sodium and it's called sodium in the names of compounds that contain this ion the third Rule now IT addresses monatomic annion they're named by taking the root of the element name and then adding ID D so the CL minus ion is called chloride right so CL is chlorine and then this an ion is now we take that root name and we add i d so chloride and we can see a couple of these examples right here here's a compound sodium chloride we have the cation sodium so we we keep it just like that as per rule two and then here we have chlorine it's the annion we take the root of the element name and then we add ID and that's why the name is sodium chloride now I want to interject right here and have a little bit of a a tangent on something important that will reappear in just a few moments but here if we're looking at sodium chloride what makes up sodium chloride we have our sodium cation and our chlorine or chloride anion how do we know that by taking sodium and chloride that we can get NAC and in this kind of ratio one sodium chloride to one one sodium to one chloride well one way to do this is by crisscrossing the charges to give you the ratio of how many cations and how many anion you need to give the chemical compound that you see here all right and the purpose of this is because for ionic compounds the overall formula of an ionic compound is charge neutral so you want to balance the charges and you can do that by crisscrossing the value that's associated in the charge with the other ion that makes up the compound so here we take this one we assign it to the CHL chloride ion and we take this one from the chloride ion assign it to sodium and so now we know we have one sodium and one chloride let's do another example of compounds and naming them and then also working through this kind of ratio and how the overall formula of an ionic compound is charged neutral how we assure that that's true when we're working through this in either direction going from the annion to the compound going from the compound to the name Etc so let's do this one this one's interesting Li i3n okay let's break this up what are the ions that make up this compound well we obviously have lithium lithium has a plus one charge and nitrogen and this nitrogen I an I has a three minus charge okay now we can go ahead and name this right off the bat we do the first part which is the cat we name the cation first and we just use the name of the element so lithium and then we can name the anion we have nitrogen but remember we take the root of the element name we add ID so it's nitride and so the name of this molecule is lithium nitride now if you were only given these ions how do you know what you would get as the chemical formula again you would crisscross these charges to tell you the ratio of the annion that you need so this one charge on lithium tells us that we need one nitrogen and the nitrogen 3 minus tells us we need three lithium so it's lithium 3n and the reason for that is if we have three lithiums with a plus one charge we have a plus three total charge from the lithium 3 and the total charge for the nitrogen is minus three and so notice how we get these charges to cancel out and remember the overall formula of the ionic compound is charge neutral now let's go ahead and do some example problems so here we're given the binary compound and we want to name each binary compound so let's look at a let's break this into the respective cation and anion we have CS all right CS this is cium all right it has a plus one charge and then we have Florine we have it as an annion F minus one all right and so it's going to be called fluoride right so we broke it into the anion it makes and we named it following the rules 2 and three for how to name cations and anion respectively and we can just put this together the name of this molecule is cium fluoride fluoride missing a you all right let's do another one all right here we have al3 Okay so let's write aluminum all right a that's aluminum it has a 3+ charge and remember cations ions are just named after the element so this is aluminum and then we have cl cl minus one here this is chlorine but for naming anion you take the root name you add ID so this is chloride and the name of this molecule is aluminum chloride beautiful let's do the last one we have lithium lithium has a plus one charge we write lithium cuz that is the cation it keeps its name and then hus this is a hydrogen anion how do we name this we take the root element name and then we add ID so this is hydride the name of this binary compound is lithium hydride wonderful so that is for naming type one binary compounds very simple rules you name the cation first then the anion monatomic cation it keeps it takes the name from the name of the element and for the monatomic annion you take the root of the element name and add i d so that's it for binary compounds type one now something else that we need to have a conversation about is well we've started with the chemical formula of a compound we learned how to decide its systematic name but the reverse process is equally important so for example if we were given the name calcium hydroxide can we write the chemical formula well calcium hydroxide gives us a lot of information calcium is ca2+ and hydroxide is O minus1 and we know that the overall formula of an ionic compound is charge neutral and to figure out how much of each ion we need we crisscross the numbers associated with charges and we figure out that the chemical formula is ca2 all right calcium forms only ca2 plus ions and hydroxide is O minus and we need two of these ions um they're required to give us a neutral compound so let's work through this even more by by doing a couple more examples cuz very important to be able to work in either direction for these kinds of problems a says pottassium iodide all right let's look at the cation cation is named first so potassium is our cation um that's k+ and then we have iodide that's our anion okay and we know iodide that is I minus all right we Criss-Cross the charges and what we get is k i so that is the formula for potassium iodide let's do this next one gallium bromide gallium is our cation it is ga3 + and then bromide is our anion it's BR minus one all right and so again we crisscross the charges because the overall formula of an ionic compound is charge neutral and since gallium has a 3 plus charge we're going to need three bromide anion to balance the charge and so our formula our chemical formula for gallium bromide is g a br3 g br3 wonderful now that we've covered that we can move into talking about binary ionic compounds type two so in binary ionic compounds type one the metal that's present forms only a single type of cation so we know sodium only forms na Plus calcium only forms form ca2+ and so on however what we're going to see now is that there are many metals specifically the transition metals that form more than one type of positive ion and thus form more than one type of ionic compound with a given annion so for example we can consider fe2 uh Fe cl2 and Fe cl3 Fe cl2 is formed with the fe2+ cation and Fe L3 is formed with the fe3+ cat in a case like this the charge on the metal ion has to be specified and so that needs to be considered in the systematic name fecl2 is actually going to be named iron 2 chloride and fecl3 would have the systematic name iron three chloride and so you notice that the Roman numerals here are used to indicate the charge of the cation and here you see a list of examples of certain metals that are that could have different charges that can form um more than one type of positive ion so we see Iron here copper Cobalt tin lead and so on Mercury you'll see this a lot again with the transition metals they're typically be able to form um more than one type of positive ion so we need to get used to being able to look at a chemical formula of a binary ionic compound type 2 and be able to give the systematic name and also be able to work in the other direction being given a systematic name for a binary ionic compound type two and be able to write the chemical formula so that's exactly what we're going to work through now so here in this first example we are we want to be able to give the systematic name for each of the following compounds here so what we're going to do is look at a first okay and let's figure out what we have for the cation and for the anion so for the cation we have copper this is one of those metals that can form multiple different positive ions okay so we need to keep that in mind while we work through this problem so copper is our cation chloride is our anion it has a minus one charge always okay now we can use this information to figure out what the charge of copper has to be here to give us cucl it's going to have to be + one so that the charge is balanced and we would get back exactly what we are looking at cucl and that helps us because now we can name it the name of this molecule is we name the cation first copper here it's a metal that can form different positive ions so we need to write the Roman numeral to indicate the charge that's associated with the copper in this formula right here and then we can name the annion is just going to be chloride wonderful let's do B now okay again same logic what do we have here we have mercury all right Mercury is one of those um elements that can have different positive it could form different positive ions okay and what is our annion we have oxide all right and the annion oxide has a charge of 2 minus this helps us determine what the charge of mercury is because the charges need to balance to give us hgo and so the charge of um Mercury here is going to be 2 plus and this is good we figured that out so now we can write the systematic name and this is going to give us if we cross charges hg2 O2 and that's okay because the reduced form of this is HG o and now we can name it it's going to be Mercury for the cation we have to use Roman numerals to indicate the charge that's going to be two and then we can now write the annion which is oxide wonderful now I want you to attempt see and let me know in the comments what you got for the systematic name of Fe2O3 this should be a three my bad so let me fix that for you it should be Fe e23 so let me know what the systematic name of this is going to be but we're going to work through this next example now together which says write the formula the chemical formula for each compound okay a tells us we have manganese for oxide this is nice though the cat is listed first manganese is our cation and we're told four so that's the charge that's associated with manganes so it's four plus and then we know that our annion is oxide oxide has a charge of 2 minus now we have to be careful here when we write the chemical formula remember the overall formula of an ionic compound is charge neutral so what we can do here is try to figure out what does that look like what we can do is we can crisscross these charges all right we can crisscross these charges it gets us mn24 but we want to reduce this down right because they are are these are factors of each other and so it's mn2 and that makes sense if I had two oxides I would have 2 minus and two of those so that's going to give me a total of four minus four minus and four plus those Balance each other out for us to get a ionic compound that is charge neutral and so the formula for manganese for oxide is mn2 let's also do this next one lead to Chloride okay so again lead is PB and it's we're giv given the charge 2+ and then we have chloride which is a minus one and then we can crisscross these charges and what we're going to get is that it's PB cl2 and that is the chemical formula for lead 2 chloride now we can move into um the next objective in in nomenclature here but before moving on I want to make note that this table has a lot of common polyatomic ions and their names this is another table that I highly recommend getting familiar with because it is super important and you're you're going to see these kinds of ions reappear throughout General chemistry and it will make General chemistry a lot easier if we're able to understand when we're given a chemical formula what is the systematic name for it or if we're given a systematic name in a problem that we know what the chemical formula for that is because as we dive deeper into General chemistry those things are going to have a huge effect on so many Concepts down the line we won't be able to do any of the other Concepts if we don't have nomenclature down so all that to say that we should familiarize ourselves with these ions so some common ones that are definitely important ammonia nh4+ we have nit nitrite that's NO2 minus and nitrate which is NO3 minus we have sulfite s so32 minus and Sul at S so4 2 minus this is hydrogen sulfate s hso4 minus again our hydroxide reappears o minus we also have cyanide CN minus we have p43 minus that's phosphate and when there's one hydrogen added it's hpo42 minus that's hydrogen phosphate when there's two hydrogens there h24 minus that's dihydrogen phosphate um here CO3 2 minus carbonate that's an important one to know hc3 minus that's hydrogen carbonate also known as bicarbonate CL minus hypochlorite a cl2 minus chlorite a cl3 minus chlorate and I should have said this as I um sometimes I can mispronounce that if I'm going too fast so again CL minus is hypochlorite and cl3 minus is hypo is chlorate um then we have cl4 minus per chlorate um this is a acetate um permanganate here M4 minus um we have di chromate chromate peroxide um so all of these really important to know so make sure that you spend some time committing these to memory now we can go ahead and move into talking about naming ionic compounds with polyatomic ions so far we've been able to deal with ionic compounds with monatomic ions it was easy to look at sodium chloride and recognize that sodium is the cation chloride is the anion and work from there but now we want to deal with ionic compounds with polyatomic ions and from this table that we've looked at there are a couple of things that I want to point out that are considered polyatomic cations and the rest are going to be polyatomic annion and being able to recognize that means now we can work with ionic compounds with polyatomic ions and easily recognize which part is the cation which part is the anion and then proceed with naming and writing chemical formulas in the same manner we have up to this point so from the table here are the polyatomic cations I'm going to write them down here polyatomic cations that we should be familiar with there's just really three main ones to keep in mind right now okay we have ammonia ammonium I should say ammonium nh4+ polyatomic cation then we have hydronium hydronium I'm not sure it's in this table but very important to know hydronium is h3o+ and then the next polyatomic cation to know is mercury 1 all right it's also the first one on this table right here okay so obviously these first two polyatomic cations Mercury one okay also written as hg2 with a 2+ charge so these are our polyatomic cations all the other things that are listed in this table are polyatomic anions and their names okay so now that we are able to distinguish polyatomic cations from polyatomic anion and we've also discussed monatomic cations and monatomic anion now we can do some more complicated problems okay so again ionic compounds with polyatomic uh ions are assigned special names that must be memorized to name the compounds that contain them and we're going to learn these by doing example problems so this first problem here says give the systematic name for each of the following compounds first one here is na2so4 let's figure out what is what the cation is and what the annion is so obviously here this is our cation na A+ this is sodium our an ion is this polyatomic ion s so4 s so4 and it's going to have a 2 minus charge hence why when we crisscross the charges the written chemical formula is na2 so4 so we use this to figure out what the charge is for the S so4 by kind of like working backwards from our crisscross so again we have sodium plus that's our cation S so4 2 minus as our annion now same rules for naming ionic compounds we keep the cation name the same as the element name and here for the annion this is a polyatomic annion let's remember what S so4 2 minus we go back to our table s so42 minus is right here that is sufate so let's write that down sulfate that means the name of this chemical formula is going to be sodium sulfate all right the name is sodium sufate awesome so nothing really changed with the naming protocol that we learned is just a matter of fact that we had to learn the names of several polyatomic ions so that we can be able to do this okay next up let's do this next one together again our cation is obvious here it is pottassium let's write that down in a different color potassium k+ that is our cation now our anion is this H2 2 po4 and it's going to have a minus charge right here okay minus charge right here we concluded that from looking at the fact that potassium there's one of them so we worked backward from our crisscross so that's our anion let's write that down H2 P4 minus let's go back to our table and see if we can figure out what that is that is right here h24 minus dihydrogen phosphate dihydrogen phosphate so we're going to write that down dihydrogen phosphate so we figured out the cation and the anion and the name of this molecule is pottassium dihydrogen phosphate go ahead and do c and let me know what you got in the comments below let's work on this next example where we want to write the formula for each so a says sodium hydrogen carbonate okay sodium is the cation that's n A+ and then hydrogen carbonate this is going to be H CO3 and you can look at the table if you need to when you're first working through these hopefully as you do more and some of these names and an anion reappear here and there it becomes a lot easier to work with so sodium hydrogen carbonate sodium is our cation na+ and then our hydrogen carbonate is H three all right and it's up here H3 minus hydrogen carbonate so let's not forget the minus and again we crisscross the charges and the formula the chemical formula for sodium hydrogen carbonate is n a CO3 wonderful let's do this next one also okay cium per chlorate okay so cium is our cat C I is going to have a + one charge and then per chlorate is going to be C4 and we can double check the charge here it's minus okay so C4 minus we crisscross the charges to get per chlorate CS cl4 so there we go that is cium per chlorate wonderful now we can move on to talking about binary compounds type three now binary calent compounds are formed between two non-metals and although these compounds do not contain ions they are named very similarly to binary ionic compounds so in the naming of binary calent compounds the following rules apply one the first element in the formula is named first using the full element name two the second element is named as if it were an anion three prefixes are used to denote the number of atoms present and four the prefix mono is never used for naming the first elements so for example C is called carbon monoxide not monocarbon monoxide and here we see some examples so n2o would be named D nitrogen because there's two nitrogens monoxide cuz there's one oxygen and the oxygen is the second element and we name second elements as if they were anion if we look at no here there is one H there's only one nitrogen but we don't say mono nitrogen monoxide we just say nitrogen monoxide and again here in this table there are the prefixes that are associated with each number so if you have one use the prefix mono two is d three is Tri four is Tetra five is Penta six is hexa seven s is HEPA 8 is OCTA 9 is Nona and 10 is DECA so that is binary compounds type three and here in this next page we kind of put it all together because we did talk about a lot and so are you dealing with a binary compound if the answer is no then the next thing you want to follow up with is well is there polyatomic ions present if the answer is no then this topic we have not covered yet we don't know how to name those kinds of um compounds just yet but if the answer is yes then we not name compounds using the procedure that is very similar to Binary compounds we identify our cation we identify our anion and then we name it as we would if we were dealing with monoatomic cations and anion but just making sure that we can recognize the nomenclature that goes with the present either polyatomic cation or polyatomic an ion are both so that is for polyatomic ions now if you are dealing with a binary compound then the next question to ask is is the metal present if no then you are dealing with binary compounds type three so binary calent compounds which is what we just talked about and you want to follow the rules we just covered there but if there are metals present then the next question to ask is well does the metal form more than one cation if the answer is no then what you have is a binary compound type one so follow those rules identify your cation identify your annion your cation you use the element name and for the annion you take the element root name and then you add ID at the end if the metal forms more than one cation then what you're dealing with is a binary compound type two here we use Roman numerals to indicate the charge that's associated with the cation so this is the workflow the flowchart that you want to fall on um whenever you are trying to work through naming compounds so with that being said let's do a mixed problem set here in the first example we're going to attempt to give the systematic name for the following compounds a is p410 so here we have two non-metals phosphorus and oxygen so this is a binary coent compound type three so prefixes are used the first element is phosphorus and there's four of them so the first part of the name is Tetra phosphorus and then we have oxygen but remember in the rules we want to name the second element as if it were an anion so oxide and there are 10 so we want to use the prefix Deca but here sometimes the a is dropped because then you're going to have like two vowels right next to each other uh typically it's written like de Cox side like this then for B we have nb205 now what we have here is a monatomic c and anion but that cation is NE oium it's a transition metal and it can have different charges so it's going to require a Roman numeral we're going to have to figure out what that is but let's write we have our cation NE oium we have oxygen this has a two minus charge here what's really cool is that yeah we do this crisscross to figure out how much of each cation and anion is this o this two from the O goes to the neoi and then if Nei had a five plus charge then it would give us back exactly this and this was a fast way of figuring that out here so that means that we have for naming we'll name the cation first NE oium charge is five and then the anion is oxide I'll let you figure out C on your own and let me know down below actually what I'll probably do is for all the ones that I've asked you to attempt I'll probably put the answers I'll try to hopefully I don't forget in the description box but later on so then you can like um look at it after you've attempted these problems to double check now let's also do this next example where we write the formula for each so we have vadium 5 fluoride so the compound contains V vadium 5 plus ions and it requires five fluoride ions for charge balance here right because vadium is going to have that five plus charge fluoride that minus one then we crisscross those charges and what we get is vf5 so that's the chemical formula for vadium 5 fluoride now for this next one we have dioxygen D fluoride so here the prefixes really help us out we have two oxygen and two fluide um Florine so O2 f to and again here I want you to attempt see on your own let me know what you got or double check your answer in the description box but with that now we can move on to our last and final category for nomenclature and that is acids an acid can be viewed as a molecule with one or more hydrogen ions attached to an annion the rules for naming acids depends on whether the anion contains oxygen so if the anion does not contain oxygen the acid is named with the prefix hydro and the suffix ick when the annion contains oxygen then the acidic name is formed from the root name of the annion with a suffix of I or o is depending on the name of the annion so here we get into those sub rules if the anion's name ends in eight then the suffix i is used and if the anion has an i ending then that is replaced with ois so here this first category acids that don't contain contain oxygens we can look at a couple of examples HF this is hydrochloric acid so notice the hydro prefix the I suffix and then we end with acid and again if we look at HCL hydrochloric acid here we have a list of acids that do contain oxygen so looking at the first one hno3 now NO3 that's the anion that this is nitrate so it ends in at that means that we want to use the suffix i and add that to the root name so the name of this acid is actually nitric acid H2 though however this is called the annion NO2 is nitrite with a i t ending and so that's going to be replaced with ois so the name of this acid hno2 is nitrous acid all right and the same goes with other examples H2 so4 s so4 is sulfate with at and so the name of this acid is sulfuric acid we use the I um suffix but if we look at H2S SO3 instead of four this ends in it and so the name of this acid is sulfurous acid with that we've completed all the topics for this chapter in the next video we'll work on problems that relate to all the topics we covered here in Chapter 2 I really hope this was helpful let me know if you have any questions comments concerns down below other than that good luck happy studying and have a beautiful beautiful day