Ib Chemistry Definitions

Front 1

What is a homogenous mixture?

Back 1

A mixture when all components are indistinguishable from each other and mix evenly with the same chemical and physical properties throughout.

Front 2

What is a heterogenous mixture?

Back 2

A mixture when the separate components are distinguishable from each other and do not mix, with distinct chemical and physical properties throughout.

Front 3

Define mole.

Back 3

The amount of a substance that contains 6.02 * 10^23 particles (Avogadro's number), the number of particles in 12g of C12

Front 4

Define molar mass.

Back 4

The mass of 1 mole of any substance.

Front 5

Define relative atomic mass.

Back 5

The weighted mean of all the isotopes of an element relative to carbon-12.

Front 6

Define empirical formula.

Back 6

The simplest whole number ratio of atoms of each element in a substance.

Ex: CH2

Front 7

Define molecular formula.

Back 7

Shows the number of atoms of each element in a molecule of the substance.

Ex: C2H4

Front 8

Define structural formula.

Back 8

Shows the arrangement of atoms and bonds within a molecule.

(drawing)

Front 9

Define relative atomic mass and how to calculate it.

Back 9

The weighted average of all the different isotopes found in nature relative to C12.

To calculate multiply the molecular weights by their respective percentages, add together and divide by 100.

Front 10

What is Avogadro's number?

Back 10

6.02 * 10^23 (12 g of carbon-12)

Front 11

Dfine solute, solvent and solution.

Back 11

The solute is dissolved into the solvent to make a solution.

Front 12

What is the molar volume of a gas?

Back 12

22.7 dm^3

Front 13

Define isotope.

Back 13

Atoms of the same element (same number of protons), but with a different number of neutrons which changes their molecular weight. Have same chemical properties but slightly different physical properties.

Front 14

Define radioactive isotope.

Back 14

Unstable isotopes where the nuclei break down spontaneously and emit raditation which can cause damage to living organisms

Front 15

Define cation and anion.

Back 15

Ion: an atom that has lost or gained an electrons and carries an electric charge

Cation: positively charged ion

Anion: negatively charged ion

Front 16

In the periodic table, what are rows called and what are columns called?

Back 16

Rows: period (across the period)

Column: group (down the group)

Front 17

Define periodicity.

Back 17

The repeating pattern of physical and chemical properties shown by the different periods.

Front 18

Define metalloid.

Back 18

An element that posseses some properties of a metal and some of a non-metal.

Front 19

What are the alkali metals?

Back 19

Metals in group 1, lithium, sodium, potassium ect

Front 20

What are the halogens?

Back 20

Elements in group 17

Front 21

What are the noble gases?

Back 21

Elements in group 18

Front 22

What groups do the s, d, p and f blocks represent/what are their names?

Back 22

s-block: first two groups

d-block: groups 3-12, transition metals

p-block: groups 13-18

f-block - separate part, lanthanoids and actinoids

Front 23

Distinguish between the metals and non-metals in the periodic table.

Back 23

p block: staircase metals on the left, non-metals on the right starting with Al

Front 24

Define atomic radius.

Back 24

The distance from the nucleus to the outermost electron, or half the distance between the nuclei of two bonded atoms of the same element

Front 25

What are the trends of atomic radius across a period and down a group?

Back 25

Down each group the outermost electron is in a higher energy level, which increases radius. Across a period the number of protons increase while electrons stay in same energy level, which pulls the atom closer together so atomic radius decreases

Front 26

Are cations/anions larger or smaller than their parent atoms?

Back 26

Cations have fewer electrons, so nucleus exhibits a greater pull on remaining electrons: smaller. Anions have more electrons so nucleus has less pull: larger.

Front 27

What are the trends of cations and anions down a group?

Back 27

They both increase in size (same as atomic radius) because electrons are added in the outer energy levels

Front 28

What is the trend for cations and anions across a period?

Back 28

Ionic radius decreases (same as atomic radius), for cations: more protons but electrons in same energy level, for anions: protons increase but electrons do not

Front 29

What are the trends of melting points down groups 1 and 17?

Back 29

group 1: melting point decreases down group because atoms become larger and strength of bonds decreases

group 17: melting point increases because the attractive forces between the molecules increase

Front 30

Define first ionization energy.

Back 30

The energy required to remove one mole of electrons from one mole of atoms (of one element) in its gaseous state.

Front 31

What is the trend of first ionization energy across a period?

Back 31

Increase irregularly across a period because electrons fill the same energy level, but increase in protons increases pull on the electrons, making them harder to remove, do not increase regularly because new sub-levels are being filled

Front 32

What is the trend of first ionization energy down a group?

Back 32

The value decreases because the outer electron is in a new energy sub-level and is therefore easier to remove

Front 33

Define electronegativity.

Back 33

A measure of the attraction that an atom has for a shared pair of electrons when it is covalently bonded to another atom.

Front 34

What is the trend of electronegativity down the group and across the period?

Back 34

Increases across a period because the # of protons increase, decrease down a group because electrons are added in (same as atomic radius)

Front 35

Define electron affinity with an equation.

Back 35

The energy chnge when an electron is added to an isolated atom in the gaseous state.

X(g) + e- = X-(g)

Front 36

List some of the shared characteristics of the alkali metals. (two equations)

Back 36

Very reactive and easily lose one electron, all undergo reactions with water to form their ions, hydroxide ions and H2 gas. They can also react with Cl, Br and I to form ionic salts

Front 37

List some of the shared characteristics of the halogens.

Back 37

Eager to gain an electron to form anions, good oxidizing agents, test for halide ions by added AgNO3, silver and halide ions form precipitate

Front 38

Define ionic bonding in a compound.

Back 38

Where the cations and anions are attracted to each other through electrostatic forces and build up a strong lattice, high melting points

Front 39

Define a polyatomic ion.

Back 39

An positive or negative ion containing more than one element

Front 40

What EN differences between two atoms define the type of bond it has?

Back 40

0-0.8 = non polar covalent

0.8-1.7 = polar colvalent

1.8+ = ionic

Front 41

Define covalent bonding.

Back 41

When atoms share one or more pairs of electrons so that each atom in the molecule achieves a noble gas configuation.

Front 42

What is the theory behind VSEPR?

Back 42

Pairs of electrons arrange themselves around the central atom so that they are as far apart as possible, greater repulsion between non-bonded electrons than bonded pairs

Front 43

What are the steps of VSEPR theory to determine the shape of a molecule?

Back 43

1. Draw Lewis Structure to determine number of electron domains

2. Work out basic shape.

3. Work out the actual shape accounting for lone pairs (Hint: for each additional lone pair, subtract 2-3 degrees from bond angle)

Front 44

What shape and bond angle will a molecule with 2 electron domains have?

Back 44

linear, 180

Front 45

List shape and bond angle for a molecule with 3 electron domains and no lone pairs

Back 45

triangular planar, 120

Front 46

List shape and bond angle for a molecule with 3 electron domains and 1 lone pair

Back 46

bent/V-shaped, 117

Front 47

List shape and bond angle for a molecule with 4 electron domains and no lone pairs

Back 47

tetrahedral, 109.5

Front 48

List shape and bond angle for a molecule with 4 electron domains and 1 lone pair

Back 48

trigonal pyramidal, 107

Front 49

List shape and bond angle for a molecule with 4 electron domains and 2 lone pairs

Back 49

bent/V-shaped, 105

Front 50

Define a resonance hybrid.

Back 50

Two different Lewis structures written for the same molecule that are correct, the true structure of the molecule is inbetween the options

Front 51

Define allotropes and list the allotropes of carbon.

Back 51

Occur when an element can exist in different forms, such as carbon where the different allotropes are diamond, graphite, graphene, fullerene

Front 52

Describe the structures of diamond, graphite, graphene and fullerene and some of their qualities

Back 52

Diamond: tetrahedal structure, Graphite: hexagonal rings, graphene: single layer of graphite, fullurene: ball of 60 C-atoms in hexagons and pentagons

Front 53

Define London forces.

Back 53

Intermolecular forces that all molecules experience, due to temperary dipoles that form when the electrons are more in one place

Front 54

Define dipole-dipole forces

Back 54

Occur in polar molecules, stronger than London

Front 55

Define hydrogen bonds

Back 55

Strongest intermolecular force, occurs in molecules with OH, FH, NH

Front 56

Define metallic bonding

Back 56

The valence electrons in metals become delocalized from the invidual atoms to create a lattice of positive ions in a sea of electrons

Front 57

Define alloy

Back 57

A compound made up of more than one metal, such as brass (copper + zinc)

Front 58

Define exothermic and endothermic reactions

Back 58

Exothermic: a reaction that releases heat

Endothermic: a reaction that absorbs heat

Front 59

Define enthalpy

Back 59

The interal energy stored in the reactions (H)

Front 60

Define standard enthalpy change of a reaction.

Back 60

The standard change in energy (delta H) measured under standard conditions (100 kPa and 298 K)

Front 61

How does a calorimeter work?

Back 61

The heat released/absorbed in a reaction is measured by the temperature increase/decrease of a known mass of water

Front 62

What is the specific heat capacity of water?

Back 62

4.18 kJ/kg/K

Front 63

What is the equation for calculating the heat energy (kJ) for a reaction? And then converting to kJ/mol?

Back 63

kJ = mass (m) * specific heat capacity * delta T

Change the sign according to endo or exo reaction, and then divide by the amount of moles

Front 64

Define Hess' law

Back 64

The enthalpy change of a reaction depends only of the difference between the enthalpy of the products and the reactions independent of the reactant pathway
detlaH1 = deltaH2 + deltaH3

Front 65

Define standard enthalpy change for formation

Back 65

The enthalpy change needed when one mole of the compoound i formed from its elements in their standard states under standard conditions

Front 66

Define average bond enthalpy

Back 66

The energy needed to break one mole of a bond in a gaseous molecule averaged over similar compounds

Front 67

Write the equation for a universal bond enthalpy

Back 67

X-Y(g) (equil) X(g) + Y(g)

Front 68

Define the rate of a reaction

Back 68

The increase in the concentration of products per unit time (mol/dm3/s)

Front 69

What is the Maxwell-Boltzmann Distribution curve?

Back 69

graph with number of particles on y axis and kinetic energy on y, shows the kinetic energy that moving particles in a gas/liquid possess, curve broadens as temp increases

Front 70

What are the 3 requiremnts of a reaction between 2 particles? (collision theory)

Back 70

1) particles must collide 2) must collide with apprpriate geometry 3) must collide with sufficient activation energy

Front 71

Define activation energy

Back 71

The minimum amount of energy required in a reaction

Front 72

Name four factors that affect the rate of reaction and how

Back 72

temperature, increases collisions bc particles move faster, increases in surface area increases the particles a reactant can come into contact with (powder), more concentrated, catalysts lower activation energy

Front 73

Define (dymanic) equilibrium

Back 73

When the forward and reverse reactions occur at the same rate which makes the concentrations of all products and reactants constant

Front 74

What is the formula for finding the reaction quotient Q?

Back 74

Front 75

When equilibrium is reached, what is Q then called?

Back 75

Kc (equilibrium constant)

Front 76

What would Kc be if the reaction

1) goes nearly to completion

2) hardly proceeds

Back 76

1) Kc is greater than 1

2) Kc is less than 1

Front 77

Define Le Chatelier's principle

Back 77

To keep Kc constant, if the concentration of one of the products/reactants is increased, the position of the equilibrium will shift

Front 78

Name 4 factors affecting the position of equilibrium and how they affect it

Back 78

Concentration: an increase in concentration of products causes the equilibrium to go the other way, Pressure: moves equilibrium towards side with less volume (particles), Temperature: increase causes endothermic reaction, Catalyst: increases rate but not Kc

Front 79

How does the ionic theory of acids and bases define an acid/base?

Back 79

Acid: substance that produces hydrogen ions in aqueous solution

Base: substance that could neutralize an acid

Front 80

Define a Brønsted-Lowry acid and base

Back 80

Acid: substance that donates a hydrogen ion

Base: substance that can accept a hydrogen ion (proton)

Front 81

Define conjugate base

Back 81

Species remaining after the acid has lost a proton, every base has a conjugate acid and so on

Front 82

Amphiprotic means

Back 82

Substance that can behave as both an acid and a base, such as water

Front 83

Write equations for the following reactions of acids: neutralization, reactive metal, carbonate, hydrogen carbonate

Back 83

acid + base = salt and water

acid + reactive metal = salt and hydrogen

acid + carbonate = salt, carbon dioxide, water

acid + hydrogencarbonate = salt, carbon dioxide, water

Front 84

Define the equation for pH

Back 84

pH = -log10[H+(aq)]

Front 85

Define the equilibrium constant for the disassociation of water

Back 85

Kw = [H+][OH-] = 1*10^-14

Front 86

Define a strong acid/base

Back 86

An acid/base that is completely dissociated into its ions in aqueous solution

Front 87

Define weak acid/base

Back 87

An acid/base that are only slightly dissociated into their ions in aqueous solution

Front 88

Define monoprotic and diprotic acid

Back 88

Monoprotic: acid that produces 1 mol of H+ in solution, HCl
Diprotic: produces 2 mol of H+ atoms in solution H2SO4

Front 89

How can you distinguish between strong and weak acids and bases?

Back 89

pH measurement, conductivity measurement (stronger will conduct more electricity), concentration measurement (rate of reaction with metals is greater with strong acids)

Front 90

Define acid deposition

Back 90

process by which acid particles, gases and precipitation leave the atmosphere, rainwater is naturally acidic so acid rain is rain with a pH of less than 5.6

Front 91

Write equations (4) and describe the deposition of sulfur oxides

Back 91

Comes from volcanoes, combustion of fossil fuels

sulfur + oxygen = sulfur dioxide

sulfur dioxde + oxygen = sulfur trioxide (sunlight)

both oxides react with water to form sulfurous acid and sulfuric acid

Front 92

Write equations (4) and describe the deposition of nitrogen oxides

Back 92

Electrical storms, bacteria, combustion and jet engines
nitrogen + oxygen = nitrogen monoxide, add more oxygen to become nitrogen dioxide
reacts with water to form nitric acid and nitrous acid

Front 93

Effects acid rain can have on the environment (4)

Back 93

Increases soil acidity which reduces important ions which can kill vegetation, decreases pH of lakes and rivers to kills fish, marble buildings become eroded, irritate mucous membranes of human lungs create resporiatory illnesses

Front 94

Three methods to counter effects of acid rain

Back 94

lower NOx and SOx produced in industry, alternative methods of energy, liming lakes by addid calcium oxide to reduce acidity

Front 95

Define oxidation and reduction

Back 95

Oxidation: the loss of electrons

Reduction: the gain of the electrons

Front 96

Describe steps for balancing a redox reaction (4)

Back 96

1) use oxidations states to define what is being oxidized/reduced

2) write half equations and balance by adding H+ or H2O

3) balance half equations for electrons

4) add together

Front 97

Define oxidizing and reducing agent

Back 97

Oxidizing agent: an agent that readily oxidizes other substances and becomes reduced

Reducing agent: reduces others and therefore becomes oxidized

Front 98

Describe the activity series of the metals and halogens

Back 98

activity series of the metals as reducing agents (become oxidized) the more reactive the metal the stronger reducing agent it is,

halogens are oxidizing agents and the stronger the halogen the more reduced it becomes

Front 99

Describe the Winkler method

(4 reactins to calculate BOD in ppm)

Back 99

magnesium + hydroxide + oxygen = manganese oxide and water

add potassium iodide, which reacts to form iodine

add sodium thiosulfate

mol ration is that 4 thiosulfate: 1 oxygen

ppm = mass in mg in 1 dm3 of water

Front 100

Define half cell and voltaic cell

Back 100

Half cell: a metal in contact with an aqueous solution of its ions

Voltaic cell: two different half cells connected by an external wire and a salt bridge

Front 101

Define electrolytic cell

Back 101

Makes a non-spontaneous reaction occur by using electrolyte (such as molten salt) and external source of electricity
electrolysis is important to obtain reactive metals from their ores

Front 102

Draw diagrams of both types of electrochemical cells.

Back 102

:3

Front 103

Define homologous series

Back 103

A series of compounds with the same general formula and functional group, successive members of a series differ by CH2, chemical properties are similar while there is a gradual change in physical properties

Front 104

Draw and write the suffix for the following compounds

Back 104

alkane, alkene, alkyne, alcohol, aldehyde, ketone, carboxylic acid, ester, amine, nitrile

Front 105

Define structural isomer

Back 105

Molecules that have the same molecular formula but different structral formulas

Front 106

Describe reactions alkanes undergo

Back 106

combustion reactions with oxygen to produce co2 and water or carbon monoxide, carbon, water (incomplete)

substitution reactions with halogen in UV light

initiation, propagation, termination

Front 107

Describe reactions alkenes undergo

Back 107

addition

polymerization

Front 108

Describe reactions alcohols undergo

Back 108

combustion

oxidation: primary alcohols are oxidized via potassium dichromate to an aldehyde then carboxylic acid

secondary are oxidized to ketones

no tertiaty oxidation