solute

part in the lesser amount

solvent

part in the greater amount

gas in gas

air

solvent = nitrogen

solute = oxygen, CO2, Ar

gas in liquid

soda

solvent = H20

solute = CO2

liquid in liquid

mixed drinks

solvent = H2O

solute = ethanol

solid in solid

alloys e.g. brass, stainless steel

solvent = zn

solute = cu

solid in liquid

salt water

solvent = H2O

solute = NaCl

mass percent/weight percent

g solute    x 100 = %w/w

g solution

ppm

(parts per million)

      g solute     

10⁶ g solution

mole fraction

            mol solute           

mol solute + mol solvent

molarity (M)

mol solute

L solution

molality (m)

mol solute

kg solvent

normality (N)

equivalents solute

L solution

M(# equivalent/mol)

Henry's Law

The solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid

concentration = k_HP

non-electrolyte

solute stays as molecules, do not ionize

e.g. glucose

all sugars

molecular compounds

electrolytes

solute dissociates in water

acids, bases and salts

metal cations + non metal anions

ionizations occurs

weak electrolytes

partial ionization (dissociation) occurs

all three aqueous species are in solution at the same time. they are in equilibrium

strong electrolytes

complete ionization (dissociation) occurs

strong acids

HCl  HBr  HI  HNO3 

HClO4  HClO3  H2SO4

strong bases

LiOH  NaOH  KOH  RbOH  CsOH  Ba(OH)2  Sr(OH)2  Ca(OH)2

weak salts

Pb(Ac)2  CdI2  HgCl2  Hg(CN)2  Fe(SCN)3

SOLUBILITY RULES

SOLUBLE

all group 1

all nitrates, acetates, chlorates, perchlorates

all chlorides, bromides, iodides except Ag, Pb, Hg2

All sulfates except Pb, Ba, Ca, Sr, Ag - slightly

INSOLUBLE

all carbonates, phosphates except group 1 and ammonium salts

all hydroxides and sulfides except group 1, ammonium salts, and Ba, Ca, Sr are slightly

n

principle quantum number (shell)

orbital size

primary energy level

increase n = increase energy

l

angular momentum quantum number

subshell

orbital shape

0=s, 1=p, 2=d, 3=f

0 < l < n-1

m_l

magnetic quantum number

orbital orientation (direction in space)

number of the same type of orbitals within the same energy (degenerate orbitals)

-l < m_l < l

m_s

spin quantum number

+1/2 or -1/2

Pauli Exclusion Principle

no two electrons in the same atom can have the same 4 quantum numbers

Aufbau Principle

for an atom in its group state you fill the lowest energy orbital first and then go up in energy until all of the electrons are used

Paramagnetism/Diamagnetism

deals with the number of unpaired electrons in an atom

more unpaired electrons = stronger magnetic field

Paramagnetism

has unpaired electrons

more unpaired e^- = more paramagnetic

diamagnetic

no unpaired electrons

Hund's Rule

maximum number of unpaired electrons, maximum amount of paramagnetism for the atom

when electrons enter degenerate (2p) orbitals, one electron goes into each orbital before 2 electrons go into any orbital

cations

lose from the highest level, not the last electron added

Newlands

1864

Law of Octaves

Every 8th element has similar properties

Meyer

1869

graphed atomic volumes vs atomic weight

Noticed: atomic volumes rise and fall periodically

alkali metals are always at the top of the curve

Mendeleev

1869

set up table with 12 rows by 8 groups

based table on: increasing atomic weight

similar properties within a group (dominant)

left spaces for undiscovered elements (Ga)

Raleigh and Ramsey

1894

isolated argon (first rare gas)

Mosely

1913

introduced concept of the atomic number

atomic number gave orderly increase with no discrepancies and atomic no longer an issue

Modern Periodic Table

current periodic table is a modified long form (expanded) Mendeleev's table with atomic numbers from Mosely

representative elements

groups 1, 2, 13-18

top inner transition period

Lanthanide

Rare Earth

bottom inner transition period

Actinide

Trans uranium

group 1

alkali metals

group 2

alkaline earth metals

group 13

boron family

group 14

carbon family

group 15

pnictogens

group 16

chalcogens

group 17

halogens

group 18

rare gases

Metals

solid at room temp (except hg)

conduct heat and electricity

malleable

ductile

shiny

non metals

non-conductive

can be solid, liquid, or gas

form anions when combined with metals

form covalent bonds amongst themselves

metalloids

appearance of metals, have non-metal properties

diagonal B to Te

shielding or screening

phenomenon where inner shell electrons block some of the positive charge of the nucleus from the outer shell electrons

amound of positive charge blocked = number of unner shell electrons

Z_eff

Effective nuclear charge

The amount of positive charge actually experienced by the outer shell electrons

Z_eff=Z - # inner shell e^- = # valence e^-

z=atomic number

radius

n²a_o

Z_eff

n=orbit #, period #

a_o= constant, leave as is

isoelectronic series

group of atoms/ions with the same number of electrons

ionization energy

minimum energy needed to remove the outermost electron from an atom in the ground state

increases across a row

decreases across a row

= -Z²_effA

n²

electron affinity

energy change when electron added to an isolated atom

largest EA = Cl

ideal gases

no attractive forces between the molecules

gas molecules have no appreciable volume

Pressure

defined as force / area

atm

1 atm = 760 mmHg

1 atm = 760 torr

Avogadro's Law

looked at changes in volume as the amount (moles) of gas changes

volume increases = n increases

V = k_An

Boyle's Law

Looked at the changes in pressure as the volume of the gas changes

pressure decreases = volume increases

PV = k_B

Charles' Law

looked at the changes in volume as the temperature of gas changes

volume increases = temperature increases

V = k_CT

Gay-Lussac

looked at the changes in pressure as the temperature of gas changes

found that the pressure increases as temperature increases

P=kT

Kelvin

K = C + 273

Avogadro's Hypothesis

STP

T = 273 K

P = 1 atm

1 mol of gas at STP = 22.4 L

combined gas laws

ideal gas law

PV=nRT

R=.0821 Latm/molK

Weight density

g/V = PM/RT

M=molar mass

Molecular weight

M= dRT / P

Dalton's Law of Partial Pressure

Mole fraction

Y

mole of a single gas

total moles of gas

Real Gases

molecules are attracted to each other

molecules take up space

observed volume > ideal volume

V_ideal=(V_obs-nb)

intramolecular

within molecule

intermolecular

between 2 separate molecules

Attractive forces effects

gas molecules are attracted to each other

this effects the force of the collisions with the container so the pressure is less than the ideal pressure

P_ideal=P_obs+((n²a)/(V²))

Van der Waals equation of real gases

combining the volume effect and the attractive forces correction

((P_obs+n²a)/(V²))(V-nb)=nRT

Kinetic molecular theory

deals with what is happening on the molecular level that causes the observed properties

POSTULATES

• gas is made up of small particles
• molecules are separated by great distances and occupy a small % of total V
• no intermolecular forces
• constant and random motion
• energy remains constant

KMT Equation

PV = (1/3)Nm(u²(avg))

N=avogadros number

u(avg)² = average of the speed squared

m=mass of molecule

Causes of pressure

P=force/area

affected by

frequency of collisions

force of collisions

Kinetic energy and KMT

KE=.5mu²

KE(avg)=1.5kT

temp increase = KE increase

depends only on temp, not gas type

molecular speed

mean free path

average distance traveled by molecules between collisions

H₂ at 0C and 1atm = 1.3x10^-5cm

Graham's law

rate of diffusion or effusion of two gases is directly proportional to the sqare root of their molecular weights

Rate A = distance travelled by A / time

Rate A/Rate B=square root of M_B/M_A

Maxwell-Boltzmann Distribution of Molecular Speeds

T increases = U_RMS increases