electrons are a fundamental particle that have a mass of 9.11 * 10 31st kg and a charge of ne1 understanding the electron is essential to understanding chemistry the behavior of an atom is determined by its relationship with its electrons as an atom will change the number of electrons it has to become more stable for example atoms can steal electrons from each other creating ions that stick together and form ions I bonds or they can share electrons creating coent bonds that fill their outer shell to help us understand how and why these bonds form we can look at how electrons behave and where they're located within the atom electrons can be found anywhere outside of an atom's nucleus but as they pop in and out of existence they're more likely to be found in some places over others we can measure the areas that electrons have been around any given atom this has created maps with defined shapes which we call electron orbitals part of the behavior of an electron comes from a property called spin this is a property with no macroscopic analog meaning the electrons are not actually spinning rather spin is used as a way to help describe an electron's angular momentum electrons can have an up or down spin which we will represent using arrows we could think of each electron orbital like a box that can hold a maximum of two electrons electrons can occupy an orbital individually or in pairs with different spins this is called the poly Exclusion Principle which we'll come back to later these orbitals are grouped into shells based on their energies which we can see using an electron orbital diagram in this diagram electron shells are labeled in order of increasing energy using a number we call the the principal quantum number n in each shell there are different types of orbitals with different energy levels these are labeled with the letters s p d or F the number of different orbital types increases with the shell number n Shell One contains only s orbitals shell two contains S and P shell three contains s p and d and so on WE indicate the shell in which the orbital is found by writing the shell number for the orbital letter for example the S orbital in Shell one is labeled as the 1s orbital the s p d and f orbitals represent the different shapes and possible energies of where electrons can exist around an atom s orbitals are spherically shaped there can only be one s orbital per shell due to its spherical nature there are no other orientations of the S orbital P orbitals are dumbbell-shaped this allows us to have three p orbitals which lie at 90° angles to each other one on the x-axis one on the Y AIS and one on the z-axis D orbitals are mostly clover-shaped like two P orbitals overlapped with each other there's one unique standout that is shaped like a p orbital with a ring around it in total there are five D orbitals F orbitals are the most uniquely shaped there are seven F orbitals in total we will see the is the least orbitals of the same type in the same shell of an atom have equal energy we call groups of equal energy orbitals subshells let's count the number of orbitals in each Shell Shell one only contains the 1s orbital shell two contains a 2s orbital and a 2p subshell with three p orbitals this gives shell two a total of four orbitals shell three contains a 3s or orbital 3 3p orbitals and 5 3D orbitals giving nine orbitals in total the pattern we see here shows that the number of orbitals within a shell is equal to the Shell number N squared the number of electrons in a neutral atom is equal to the atom's atomic number these electrons will fill an atoms orbitals in subshells in a very particular way the first of three guidelines that we will follow is called the B principle this principle states that electrons enter subshells in order of increasing energy meaning the 1s orbital is filled out first followed by the 2s orbital then the 2p subshell then the 3s orbital then the 3p subshell and so on we saw our next guideline a bit earlier this was the poly Exclusion Principle which again states that orbitals can hold no more than two electrons which must have opposite spin we see this in our electron orbital diagram where orbitals with two electrons are shown as boxes with an up and down arrow take this orbital in 3p for example while electrons can enter the same orbital if they have opposite spin this is not usually the most stable configuration consider electrons filling into the 3p sublevel when two electrons are paired within the same orbital they repel each other electrostatically Huns rule our final guideline stat that due to this repulsion electrons are placed in separate orbitals of a subshell if possible within a subshell we would also like to maximize the spin of our electrons it's more stable for an atom to contain electrons with the same spin as opposite spin electrons even in different orbitals repel each other more strongly therefore when we add electrons to a set of orbitals within a subshell we will add the electrons to one orbital at a time each with the same spin when each orbital has one electon we will then begin doubling up with electrons of opposite spin the arrangement of electrons in an atom is called its electron configuration this will show electron shells and subshells in order of increasing number in energy with a superscript indicating the number of occupying electrons in each subshell for example a nitrogen atom seven electrons have an electron configuration of 1 S2 2 S2 2 B3 we can see this in reference to the electron orbital diagram and the atoms bore model while the electron orbital diagram matches identically to the configuration we lose some information about our electrons when we solely rely on a bore model so while this is a nice visual the model is limited and doesn't tell us the whole story what happens when an atom changes their number of electrons when atoms gain or lose electrons they form ions with positive or negative charges these electrons are added or removed from an atom's highest energy orbital for example oxygen has eight electrons with an electron configuration of 1 S2 2 S2 2p4 when oxygen forms the oxide ion it will gain two electrons forming a charge of 2 minus it will have a new electron configuration of 1 S2 2 S2 2 P6 with the two extra electrons added to the 2 p suev for atoms with many subshells it could become too timec consuming to write out their full electron configuration to get around this we can recognize that the inner shells of an atom might share the electron configuration of a neutral noble gas which have full shells of electrons take for example phosphorus and its electron configuration compared to its nearest noble gas of neon we can see that the inner electron shells of phosphorus line up exactly with the full electron configuration of neon we can therefore write the noble gas configuration of phosphorus as neon in square brackets followed by 3s2 3p3 some subshells have similar energies despite being in different shells take for example 3D and 4S although the 4S orbital is in a further shell it will usually fill before 3D this is unusual to our typical pattern to remember how electrons fill we can draw a triangular version of our electron orbital diagram by tracing diagonally through our diagram we could track The Filling order of an electron configuration Let's test this with iron iron has the electron configuration of 1 S2 2 S2 2 P6 3 S2 3 P6 4s2 3d6 when electrons are removed from an atom they're taken from the the highest energy subshell first in a 3D metal such as iron when the 4S and 3D subshells are occupied with electrons the order of their energies switch when iron forms the iron 3+ ion for example electrons are removed from 4S and then 3D this would change iron's electron configuration to end in 3s2 3p6 3d5 with the 4S orbital now empty there are a couple more exceptions for how electrons fill electron orbitals one outcome of Hun's rule in electron pair repulsion is that sometimes it's more stable for 3D and 4S orbitals to either be all singly occupied or completely filled this is the case for chromium typical electron configuration rules would give an electron configuration ending in 4s2 3d4medical.com figuration actually ending in 4s1 3d5 conversely we have copper typical guidelines would Place Copper's outer electrons in 4s2 3d9 however it's been found that Copper's 3D subshell is fully occupied rather than 4S its electron configuration actually ends in 4s1 3d10 in summary the behavior of electrons forms the basis of most chemistry electrons dictate how atoms interact and how chemical reactions occur the different ways of representing electrons helps us understand different aspects of their behavior or models for example give us an easy visual for electron shells in the size of atoms whereas electron probability clouds show where electrons most likely exist and electron configurations give us an address for the orbitals and subshells that electrons fill using each of these models we can better understand the foundation of how atoms bond and react with each other