General Chemistry Master Set

element

cannot be broken down chemically into simpler substance

compound

• two or more elements chemically joined in fixed ratio

• has different properties from elements it formed from

mixture

• two or more elements not chemically bonded together

• retains properties from elements it formed from

solid to gas

sublimation

gas to solid

deposition

solid

has fixed shape and volume

liquid

has fixed volume but no fixed shape

gas

has neither fixed shape or volume

physical change

• no new substances are produced

• melting of ice is an example

chemical change

results in the formation of new chemical substances

atomic structure

• protons and neutrons (nucleons) are located in nucleus of atom

• electrons are located in energy levels surrounding nucleus

principal energy level (n)

• electrons are located in principal energy levels

• the first energy level (n=1) has lowest energy, and energy increases as value of n increases

• each main energy level can hold at most 2n² electrons

• the atomic number (Z) is the number of protons in an atom

• the mass number (A) is the number of protons and neutrons (nucleons) in an atom

• atoms have same number of protons as electrons

positive ion

less electron than proton

negative ion

more electron that proton

isotope

• elements that have same atomic number but different mass number

• as isotope mass increases, boiling point, melting point, and density increases

relative abundance

percentage of atoms with a specific mass number in naturally occurring environment

relative atomic mass

• mean of the atomic mass of each isotope weighted by relative intensities

• only use relative intensities for singular element or diatomic element (but not both)

electromagnetic spectrum

• when there is short wavelength, there is high frequency and energy

• when there is long wavelength, there is low frequency and energy

• violet, indigo, blue, green, yellow, orange, red are in order of shortest to longest wavelength

continuous spectrum

shows all wavelengths of visible light

absorption line spectrum

black lines on coloured background

emission line spectrum

• coloured lines on black background

• spectral lines converge at higher energy or shorter wavelength

spectroscope

splits light into different wavelengths

bohr model

• when electrons absorb photon of energy they transition to higher energy level

• when electrons transition to lower energy levels they emit photons of energy

energy of photon emitted for electron transition

E=hf

• f is frequency of emission spectrum created by photon

complete emission spectrum

• electron transitions to n=3 emit infrared radiation

• electron transitions to n=2 emit visible light

• electron transitions to n=1 emit ultraviolet light

atomic orbitals are regions where there is a high probability of finding an electron

top is s orbitals

bottom is p orbitals

sublevels

• each principal energy level is split up into sublevels

• n=1 has 1 sub level (1s)

• n=2 has 2 sub levels (2s, 2p)

• n=3 has 3 sub levels (3s, 3p, 3d)

• n=4 has 4 sub levels (4s, 4p, 4d, 4f)

• within a principal energy level the order of energy is s<p<d<f

• s, p, d, f hold 2, 6, 10, 14 electrons respectively and have 1, 3, 5, 7 orbitals respectively

aufbau principle

lowest energy sublevels are filled first

pauli exclusion principle

atomic orbital must have maximum of two electrons and have opposite spins

hund’s rule

orbitals in a sublevel are filled with one electron each with same spin before being doubly filled

positive ion configuration

lose electrons from highest principal energy level

negative ion configuration

add electrons based on aufbau principle

ions

have same electronic configuration as noble gases

convergence limit

• the principal energy levels converge at higher energy

• the spectral lines also converge at higher energy

ionisation energy

• energy required for the electron transition from n=1 to n=infinity

• use E=hf, where f is frequency of convergence limit to get ionisation energy of 1 atom. multiply that by avogadro’s number to get energy per mol.

trend in successive ionisation energy

• ionisation energy increases as more electrons are removed from gaseous atom due to greater attraction from more positive ion

• large increases in ionisation energy when maximum principal energy level changes

ideal gas assumptions

• particles in ideal gas are in constant, random, straight line motion

• collisions between particles of an ideal gas are elastic

• the volume occupied by the particles of the gas are negligible

• there are no intermolecular forces between the particles of an ideal gas

when is gas closest to ideal

• gases are most ideal at high temperatures and low pressures

• they deviate the most at low temperatures and high pressures

cation

positive ion

anion

negative ion

polyatomics

sulfate: SO4 2-

sulfate: SO3 2-

phosphate: PO3 3-

nitrate: NO3 -

nitrite: NO2 -

carbonate: CO3 2-

ammonium: NH4 +

hydroxide: OH -

• atoms in polyatomic ion are covalently bonded

• bonds in compound that contains polyatomic are ionic

ionic compound properties

• have a lattice structure

• solids under standard conditions

• high melting and boiling points

• not conductive when solid, but conductive when molten or dissolved

• soluble on polar solvents (like water) where compound is split up into ions

• ions are surrounded by water molecules (hydration)

factors impacting melting point of ionic compound

• greater the charges on the ions, the higher the melting point

• smaller the ionic radius the higher the melting point

lattice enthalpy

• energy change when one mol of solid ionic compound is broken down into gaseous ions

• endothermic process with a positive enthalpy change

factors effecting lattice enthalpy

• greater the charges on the ions, the higher the lattice enthalpy

• smaller the ionic radius the higher the lattice enthalpy

nobel gases

group 18

octet rule

• atoms bond together to achieve full valence shell of 8 electrons

• nobel gases have full valence shells so they are very stable

ionic bonding

• electrostatic attraction between oppositely charged ions

• always forms between metal (cation) and nonmetal (anion)

• transfer electrons to achieve full octet

covalent bonding

• share electrons to achieve full octet

• occurs between two non metal elements

octet rule exceptions

• hydrogen, beryllium, and boron are stable with less than 8 valence electrons

• period 3 elements are stable with more than 8 valence electrons (expanded octet)

multiple covalent bonds

• in single, double, and triple covalent bond there are 2, 4, 6 shared electrons respectively

• bond energy increases and bond length decreases with number of covalent bonds

coordinate covalent bond

one atom contributes both electrons in a bond

VSEPR theory

• bonds repel each other to be as far as possible

• lone pairs repel each other to be as far as possible

electron domain

can be lone pair, or single/double/triple bond

molecular shapes

electronegativity difference bonding

• 0-0.4 non polar covalent

• 0.5-1.7 polar covalent

• >= 1.8 ionic

bond dipole

• partial negative charge on atoms with high electronegativity

• this is due to more electronegative atom pulling shared electrons closer

arrow points towards more electronegative atom

non polar molecules

• linear, trigonal planar, and tetrahedral molecules with same atoms bonded to the central atom

• diatomic molecules with same atoms bonded

• molecules with very weak polar bonds (C-H)

polar molecules

• linear, trigonal planar, and tetrahedral molecules with different atoms bonded to the central atom

• diatomic molecules with different atoms bonded

• v-shaped molecules with polar bonds

molecular covalent

• covalent substances that exist as individual molecules

• intermolecular forces between those molecules

• low melting and boiling points due to weak intermolecular forces

• polar molecules are soluble in polar substances, nonpolar molecules are soluble in nonpolar substances

• nonconductive

giant covalent

• covalent bonds that form a lattice structure

• there are no intermolecular forces

• high melting and boiling points due to strong covalent bonds

• insoluble and nonconductive

allotrope

different forms of the same element

graphite allotrope of carbon

• trigonal planar geometry

• weak intermolecular forces between layers so layers can slide over each other

• good conductor because of delocalised electrons

diamond allotrope of carbon

• giant covalent structure

• high melting and boiling point

• tetrahedral geometry

• nonconductive

fullerene C60 allotrope of carbon

• 12 pentagons and 20 hexagons

• trigonal planar geometry

• some conductivity but not as much as graphene

graphene

• one layer of graphite (very thin)

• good electrical and thermal conductivity

LDF forces

• instantaneous dipole formed by random movement of electrons

• as molar mass increases strength of LDF increases, which results in higher boiling point

dipole-dipole forces

• dipole-dipole forces occur between two polar molecules between partial positive portion of one molecule and partial negative portion of another

• stronger the dipole moment the higher the boiling point

hydrogen bonding

• very strong dipole-dipole forces caused by hydrogen and F/O/N bond

• hydrogen bonding is reason for water’s high boiling point compared to other liquids

strength order of forces

LDF < dipole-dipole < hydrogen bonding < covalent bonding < ionic bonding

factors affecting solubility

• polar molecules are soluble in polar substances, nonpolar molecules are soluble in nonpolar substances

• most ionic compounds are soluble in water due to strong polar nature causing ion-dipole forces

factors affecting conductivity

• covalent substances that exist as molecules have poor conductivity because their electrons are localised to be near to molecule

• giant covalent structures also have poor conductivity because their electrons are localised in covalent bonds

• when ionic compounds are melted or dissolved their ions are free to move and conduct electricity

• metallic substances are good conductors because there is sea of delocalised electrons that are free to move

thin layer chromatography

• used to seperate substances based on their solubility in water

• rf value is distance substance travels over distance travelled by solvent

• more soluble substances have higher rf values

• these can be compared to determine what substances were separated

sigma bond

• formed along internuclear axis

• between two s orbitals/s and p orbital/2 p orbitals

pi bond

• formed over internuclear axis

• formed between two p orbitals

sigma and pi bonding in covalent

• single covalent: 1 sigma bond

• double covalent: 1 sigma and 1 pi bond

• triple covalent: 1 sigma and 2 pi bond

delocalised pi electrons

• exist in all molecules with for which there are resonance structures

• electrons are free to move across ions as the orbitals combined

hybridisation

• mixing of orbitals to produce hybrid orbitals used for covalent bonding

• sp3 hybridisation has 4 orbitals (one s and three p), occurs with 4 electron domains

• sp2 hybridisation has 3 orbitals (one s and two p) and leaves 1 unhybridised p orbital, occurs with 3 electron domains

• sp hybridisation has 2 orbitals (one s and one p) and leaves 2 unhybridised p orbital, occurs with 2 electron domains

resonance structure

• occurs when double bond can be in multiple places in compound

• the actual bonds are intermediate between single and double bond

formal charge

• used to determine preferred lewis structure of compound

• preferred lewis structure has near 0 formal charge for all atoms in compound, negative charges on electronegative atoms

• FC = V (valence) - N (non bonding) - B (bonding) / 2

metallic bonding

electrostatic force between positive lattice and sea of delocalised electrons

factors effecting metallic bonding strength and melting point

• strength of metallic bond increases with greater ionic charge and smaller ionic radius

• this also results in a higher melting point

properties of metallic structures

• have nondirectional bonds

• good conductors of heat and electricity due to presence of delocalised electrons

• malleable (can be bent into shape because of layers easily shifting)

• ductile (can be drawn into wires)

alloy

alloys are materials that are composed of two or more metals

properties of alloys

alloys are less malleable (harder) because different size atoms make it harder for layers to shift

position of metals, nonmetals, and metalloids on periodic table

properties of metalloids

• have properties of both metals and nonmetals

• solids at room temperature

• some are shiny

• brittle

• intermediate electrical conductivity

• moderate density

properties of nonmetals

• gases at room temperature

• dull

• brittle

• poor electrical conductivity

• low density

• high ionisation energy

• high electronegativity

properties of metals

• solids at room temperature

• shiny

• malleable and ductile

• high electrical conductivity

• high density

• low ionisation energy

• low electronegativity

elemental blocks

• block indicates what sublevel electrons are being added to

• d block elements are also known as transitional elements
  

electron shielding

inner electrons repel the valence electrons, shielding the valence electrons from the attractive force of the nucleus

trends in electron shielding

• electron shielding is constant across a period

• electron shielding increases down a group

effective nuclear charge

• the effective charge experienced by valence electrons

• it is the nuclear charge minus the charge of the shielding electrons

• effective charge is constant across a period

• effective charge increases down a group

• (basically same trend as electron shielding)

trends in atomic radius

• atomic radius decreases across a period

• atomic radius increases down a group

trend in ionic radius across period 3

• ionic radius decreases across period 3 for ions of Na, Mg, Al, Si

• there is a large jump in ionic radius at P, followed by a decrease for ions of S, and Cl

• positive ions have smaller ionic radius than original atom, and negative ions have larger ionic radius

first ionisation energy

energy required to remove one mole of electrons from one mole of gaseous neutral atoms

trend in 1st ionisation energy

• across a period the first ionisation energy increases, as the number of protons increases, which decreases ionic radius, and increases attraction

• large jump down when going up a principal energy level

• discrepancy when electron is removed from new subshell, as it easier to remove electron from more energetic subshell

• discrepancy when electron is removed from subshell which has 1 pair, as it easier to remove electron that is repelled by another one

electron affinity

• 1st electron affinity is the energy released when one mole of electrons is added to one mol of neutral gaseous atoms (negative)

• 2nd electron affinity is positive due to extra repulsion when trying to add an electron

trends in 1st electron affinity

• greater electron shielding and greater ionic radius reduces the energy released

• energy released decreases down a group

• energy released is more for nonmetals than for metals