Unit 1:Atomic Structure Part 2 CH1120

Subatomic Particles Atoms are composed of Protons Electrons Neutrons The amount of these subatomic particles alters the identity of the atom

Subatomic Particles

Protons Positively charged Relative mass of 1 Found in nucleus

Neutrons Neutral Relative mass of 1 Found in nucleus

Electrons Negatively charged Negligible mass Found outside of the nucleus

Subatomic Particles When an atom has an equal number of protons and electrons it is neutral Has no charge

Ions When an atom has more protons than electrons it is a cation Positively charged, name is same as element When an atom has more electrons than protons it is an anion Negatively charged, name ends in –ide (ie. Chloride) An atom with a charge is called an ion

Ions Atoms gain and lose electrons in interactions to have the same number of electrons as the closest noble gas Na+ loses 1 electron to become like Ne S2- gains 2 electrons to become like Ar Carbon can gain or lose 4 electrons to become like Ne or He

Atomic Number Number of protons in the nucleus of an atom Found on the periodic table

Mass Number Total number of protons and neutrons in the nucleus of an atom Mass number is always larger than atomic number

Isotopes Isotopes are atoms of the same element that have varying masses Remember we can change the number of neutrons without changing the element This simply changes the mass of the atom

Isotopes Changing number of protons changes the element Atomic number Changing number of neutrons changes the mass Creates different isotope Changing the number of electrons changes the charge

Isotopes

Isotopes

Isotopes When given names like Carbon-13, 13 is the mass number Neutrons and protons

Example Draw the isotope symbol of Hydrogen-3 with a charge of +1. How many of each of the subatomic particles are present?

Isotopes  

Isotopes  

Isotopes Smoke detectors Archaeological dating Medical imaging Americium-241 Archaeological dating Carbon-14 Medical imaging Medical treatment

Average Atomic Mass Atomic mass and mass number are NOT the same thing Mass number = number of protons and neutrons Whole number Atomic mass = average atomic mass/weight

Average Atomic Mass The average of all of the isotopes of a particular element found in nature This can be found on the periodic table or can be calculated from data given

Average Atomic Mass Average Atomic Mass = [M1 x %Abundance1] + [M2 x %Abundance2] + ...

Average Atomic Mass Gallium has two naturally occurring isotopes: Ga-69 with mass 68.9256 amu and abundance of 60.11% and Ga-71 with mass 70.9247 amu and abundance of 39.89%. Calculate the atomic mass of gallium.

Average Atomic Mass Bromine has two naturally occurring isotopes (Br-79 and Br-81) and an average atomic mass of 79.904 amu. If the mass of Br-81 is 80.9163 amu, what is the mass of Br-79? The natural abundance of Br-79 is 50.69%.

Atomic Orbitals Electrons of an atom are attracted to the positive nucleus of the atom Electrons are not all the same distance from the nucleus Therefore they have different energies Further away = greater energy

Quantum Mechanical Model Electrons can be grouped according to their ease of removal from the atom Ease of removal of electrons depends on their distance from the nucleus

Quantum Mechanical Model

Quantum Mechanical Model Electrons are found in energy levels called shells Each shell is broken down into subshells Each subshells has one or more atomic orbitals with a specific 3D shape and energy Think of orbitals as clouds where electrons live

Quantum Mechanical Model

Quantum Mechanical Model F Orbitals

Quantum Numbers Principle Quantum Number (n) Identifies energy level (therefore number of subshells) Positive integral numbers (1 and up) As “n” increases, orbital size increases, electron energy increases

Quantum Numbers Angular momentum quantum number (l ) Defines the shape of the orbital All value 0 to n-1 Letters are used to represent different values of l 0 = s 1 = p 2 = d 3 = f

Quantum Numbers Magnetic quantum number (ml ) Describes 3D orientation of orbital in space Values from –l to +l

Quantum Numbers Spin quantum number (ms) Identifies the rotation of the electron No 2 electrons can have all 4 quantum numbers the same Values are +1/2 and -1/2

Quantum Numbers

Quantum Mechanical Model Electron shell: collection of orbitals with the same principle quantum number (n) 3s, 3p, 3d, and 3f, are all in the 3rd electron shell Subshell: collection of orbitals within a shell 3d is a subshell of the 3rd electron shell

Quantum Mechanical Model The Principle Quantum Number (n) determines how many subshells are present n = 1 has one subshell, n = 2 has two subshells, etc. Each subshell consists of a specific number of orbitals s = 1, p = 3, d = 5; f = 7 The total number of orbitals in a given shell is n2 Ex. n = 1 has one orbital, n = 2 has four orbitals, etc.

Orbital Shapes s-orbitals Begin at the first energy level Spherical Hold a maximum of 2 electrons

Orbital Shapes p-orbitals Begin at the second energy level 3 p-orbitals at each energy level Each orbital holds 2 electrons (total of 6 at each energy level)

Orbitals Shapes d-orbitals Begin at third energy level 5 d-orbitals at each energy level Each orbital holds 2 electrons (total of 10 at each energy level)

Orbital Shapes f-orbitals Begin at the fourth energy level 7 f-orbitals at each energy level Each orbital hols 2 electrons (total of 14 at each energy level)

Filling Orbitals When filling orbitals, electrons fill those with the lowest energy first Orbitals with the same energy are said to be degenerate Ex. All orbitals in the 3p subshell are degenerate Repulsions between electrons cause the subshells in a given shell to be at different energies

Filling Orbitals Generally, a higher n indicates a higher energy and s<p<d<f (s having the lowest and f the highest energy) But when you move past 3p, things get more complicated

Aufbau Principle Electrons will assume the most stable position based on the nucleus of an atom and the electrons already present

Pauli Exclusion Principle No two electrons can have the same exact 4 quantum numbers Orbitals can hold a maximum of 2 electrons therefore they must have opposite spins

Electron Configurations Each component consists of: A number denoting the energy level A letter denoting the type of orbital A superscript denoting the number of electrons in those orbitals

Electron Configurations Remember each orbital holds a maximum of 2 electrons s-orbitals have 1 subshell p-orbitals have 3 subshells d-orbitals have 5 subshells f-orbitals have 7 subshells

Electron Configurations 1. Determine total number of electrons for atom 2. Start filling in orbitals according to the order of increasing energy (we will look at 2 methods) 3. Continue until all electrons are used up in the orbitals

Electron Configurations

Electron Configurations

Electron Configurations Write the electron configurations for the following elements: Oxygen Iron

Electron Configurations Write the electron configurations for the following elements: H B Cl

Orbital Diagrams Each box or line represents one orbital Half-arrows represent the electrons The direction of the arrow represents the relative spin of the electron Example: lithium

Hund’s Rule For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized Example: oxygen

Orbital Diagrams 1. Determine total number of electrons for atom 2. Start filling in orbitals according to the order of increasing energy 3. Apply Hund’s rule for the subshells until all electrons are used up in the orbitals Place 1 electron in each orbital until all contain 1 electron, then start pairing

Orbital Diagrams Draw orbital diagrams for the following elements: Be Na

Condensed Electron Configurations Write the nearest noble gas with lower atomic number in square brackets Begin the electron configuration from that point

Condensed Electron Configurations Write the condensed electron configuration for Fluorine.

Condensed Electron Configurations Write condensed electron configurations for the following elements: Beryllium Chlorine Cobalt

Valence Electrons Electrons filling the outermost energy levels Used for chemical bonding Group A elements, the group number is the number of valence electrons

Valence Electrons How many valence electrons are in the following elements: Fluorine Magnesium Carbon

Valence Electrons