Unit 1: Atomic Structure and Properties

molecular weight

the sum of the atomic weights of the atoms in the chemical formula of the substance

amu

atomic mass unit

mole

a unit of measurement that is the amount of a pure substance containing the same number of chemical units, weighs the same number of grams as one of those particles weighs in amu’s

hydrate

compound that attracts and absorbs water

molar mass

used as a conversion factor to convert between moles and grams

avogadro’s number

6.02 × 10²³ particles/mole is the conversion factor to convert between the number of particles and moles

formula unit

a set of an ionic compound

information mass spectrometer provides

1. the number of isotopes present

2. atomic mass of each isotope

3. the relative amount of each isotope

average atomic mass

(relative abundance of isotope n (y-axis)) x (mass of isotope n (x-axis))

pure substance

a substance with constant composition, can be either an element or a compound

law of definite proportions

compounds with the same elements in the same proportion are the same compound

percent composition

the percent by mass of each element that makes up a compound, calculate by dividing the mass of each element in a compound by the total molar mass of the substance

empirical formula

represents the simplest ratio of one element to another in a compound

molecular formula

represents that actual formula for the substance

steps to determine empirical formula (when given element percentages)

1. assume sample is 100g to change percents to grams

2. take mass of each element and get moles of each

3. divide each mole value by lowest of the values

4. round to whole number if within 0.1, if not multiply by a factor that gives all whole numbers

5. take values and use as subscripts for each element

steps to find molecular formula (when given molar mass of substance)

1. find mass of empirical formula

2. divide molar mass by empirical formula mass to find the whole number factor

3. multiply all subscripts in the empirical formula by the value

anhydrate

substance without water

importance of heating hydrates multiple times

ensures all water has been removed from the substance

combustion analysis

burning of a hydrocarbon to determine its empirical formula

mixture

when two or more pure substances (elements and compounds) are combined, used to determine mass composition of each substance in this

stoichiometry

the determination of the proportions in which elements or compounds react with one another

mass percentage

divide the mass of the substance by the total mass of the mixture and multiply by 100

elemental analysis

part of analytical chemistry, used to determine the composition of a mixture and can be qualitative or quantitative

steps of precipitation gravimetry

1. analyte is converted to an insoluble precipitate

2. precipitate is filtered, washed free of impurities, dried completely, and weighed

3. determine the amount of the analyte in the sample (percent composition)

analyte

substance that is being studied

limiting reactant

a reactant that is totally consumed when the chemical reaction is completed

precipitate

a solid formed by a change in a solution, often due to a chemical reaction or change in temperature that decreases solubility of a solid

nucleons

another name for protons and neutrons

isotope

atoms with the same number of protons but different numbers of neutrons

steps to figure out the number of neutrons in most common isotope

1. round mass number to nearest whole number

2. subtract the number of protons to get the number of neutrons

protons and neutrons

make up the majority of atom’s mass

electrons

make up the majority of atom’s volume

electron configurations

distribution of electrons of an atom or molecule in atomic or molecular orbitals using superscripts to represent electrons

orbital notation

a way of writing an electron configuration to provide more specific information about the electrons in an atom of an element using arrows to represent electrons

core electrons

all electrons not in the very outer electron shell

valence electrons

electrons in the very outer electron shell

spdf

letters of energy levels

sphere

s-orbital shape

dumbbell

p-orbital shape

clover

d-orbital shape

tetrahedral

f-orbital shape

first shell

only has room for the 1s orbital

second shell

has room for a 2s orbital and a set of three 2p orbitals

third shell

had room for a 3s orbital, a set of three 3p orbitals and a set of five 3d orbitals

2 electrons

number of electrons each shape can hold, but there can be multiple of each shape

Aufbau principle

electrons are added to the lowest orbital first and build up

Hund’s rule

each orbital should have one electron before any are doubled up

subshell

another word for orbital

Pauli exclusion principle

no two electrons can have the same set of 4 quantum numbers (spinning in opposite directions)

noble gas abbreviation

abbreviated form of electron configurations that uses the noble gases of the periodic table

ground state

the lowest allowed energy state of an atom, molecule, or ion; the most stable configuration

excited state

any state with energy greater than the ground state that cause electrons to reach higher orbital levels and fall back down and give off color

electromagnetic forces

governs chemistry, a type of physical interaction that occurs between electrically charged particles

factors of electromagnetic force strength

1. amount of charge (higher charges have stronger attractions)

2. distance between charges (the closer together the oppositely charged things are, the stronger the attractions)

Coulomb’s law formula

F=k q₁q₂/r²

Coulomb’s law

1. force between charged particles is proportional to the product of the two charges and the force is inversely proportional to the squared radius between them

2. forces will decrease the further away the particles are

3. higher charges and smaller distances between the charges result in a greater force of attraction. this explains why it takes more energy to remove electrons that are closest to the nucleus

photoelectron spectroscopy

experimental technique that measures the relative energies of electrons in atoms or molecules

photoionization

works by ejecting electrons from the materials using high energy electromagnetic radiation and then measuring the kinetic energy of those electrons

pes graphs

show the relative number of electrons and their corresponding binding energy

binding energy

amount of energy needed to remove an electron from an atom

coulombic attraction

1. explains most periodic trends

2. negative electrons in the electron cloud and positive protons in the nucleus are attracted to each other

3. the larger the charge, the more attractive forces between the particles

4. more protons=stronger

5. further away particles are from each other, the weaker the attractive forces

6. more shells=weaker

pes graph origin

represents the nucleus of an atom

same valence electrons configuration

means they tend to have similar chemical properties

periodic trends

can be explained by the arrangement of the electrons and the number of protons in the atoms

first ionization energy

energy required to remove the outermost (highest energy) electron from a ground state, neutral atom in its gaseous from

highest first IE

element with the least shells because electrons are closer to the nucleus, which means the attraction is stronger and requires more energy

next highest first IE

elements with the same number of shells and more protons because more protons attract more electrons and the attraction will be stronger so it will require more energy

subsequent ionization energies

energy required to remove the second, third, and so on electron from a ground state atom in its gaseous form

stable element

causes a large jump in its ionization energy

atomic radius

a measure of the size of an ion

larger radius

caused by more shells because there is a larger electron cloud

smaller radius

caused by the same number of shells but more protons because the protons pull the valence shell tighter

cations

positive ions that are always smaller than the parent atom because there are fewer electron-electron repulsions

anions

negative ions that are always larger than the parent atom because electrons are added to the same valence shell, however, there are greater electron-electron repulsions so the ion increases in size

electron affinity

energy change that occurs when an electron is added to a gaseous atom or ion

electron affinity forrmula

E_(g) + e^- —> E^-_(g)

electron affinity trends

1. electron affinity increases from left to right on the periodic table because electrons are filling the valence shell and effective nuclear charge (therefore coulombic attraction) is increasing

2. in group 1 and in general, electron affinity decreases down a group because the distance from the nucleus to the valence shell increases (decreasing coulombic attraction)

electronegativity

measure of the ability of an atom (or group of atoms) to attract shared electrons in a bond

weaker electronegativity

caused by more shells because the nucleus is farther from the valence shell

stronger electronegativity

caused by same shells but more protons because more protons attract valence electrons tighter

fluorine

most electronegative element

noble gases

don’t have electronegativity values because they typically do not from bonds since they are already stable

ionic bond

always involves the transfer of electrons from the least electronegative species to the most electronegative and are traditionally described as being between a metal and a nonmetal

ionic bond element types

1. cation: element that loses electrons becomes positive

2. anion: element that gains electrons becomes negative

outer valence shell

elements lose or gain electrons to fill (s²p⁶)

electrostatic force

force that holds ionic compounds together (coulombic attraction)

nonmetals

gain electrons to filler their octet

metals

lose electrons to have a pseudo-noble gas configuration

silver ion

Ag⁺

cadmium ion

Cd²⁺

zinc ion

Zn²⁺

ionic compounds

number of electrons lost by the metal must equal the number of electrons gained by the non-metal when these are formed