As Chemistry Definitions And Need-To-Knows (01 Spec)

Front 1

Define: atomic number.

Back 1

The atomic (proton number) is the number of protons in the nucleus.

example:

12 C
6 - proton #

Front 2

Define: mass number.

Back 2

The mass (nucleon) number is the total number of protons and neutrons.

example:

12 - mass #
6 C

Front 3

Define: First Ionisation energy.

Back 3

The first ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of ions wih a single positive charge.

example:

X(g) -> X+(g) + e-

Front 4

Define: Mole.

Back 4

One mole is the amount of substance that contains as many particles as there are atoms in exactly 12g of carbon-12.

The number of particles in one mole is 6x10^23 (Avogadro's constant.)

Note: the relative atomic mass in grams of any element contains one mole of atoms.

Front 5

Define: Relative molecular mass.

Back 5

The relative molecular mass of a compound is the mass of its fomula unit - relative to 1/12 the mass of an atom of carbon-12.

Front 6

Define: theoretical yield.

Back 6

Theoretical yield is the maximum amount of product that can be made from the available reactancts. (This is calculated from the amounts of the reactants using the balanced equation.)

percentage yield = ( actual yield / theorectical yield ) x 100

Front 7

Define: Actual yield.

Back 7

The Actual Yield is the amount of product that is actually made/measured at the end of the reaction.

percentage yield = ( actual yield / theorectical yield ) x 100

Front 8

Define: first electron affinity.

Back 8

The first electron affinity is defined as the energy change when one mole of gaseous atoms accepts one mole of electrons to form one mole of ions with a single negative charge.

X(g) + e- -> X-(g)

Front 9

Define: electronegativity.

Back 9

Electronegivity is the power of an atom to attract electrons towards itself in a covalent bond.

Front 10

Define: molecular orbital.

Back 10

An orbital is the area of space where an electron is most likely to be found.

Front 11

Define: the standard enthalpy change of combustion.

Back 11

∆H⊖c
The standard enthalpy change of combustion is the enthalpy change when one mole of a substance burns completely in oxygen under standard conditions.

Front 12

Define: the standard enthalphy change of formation.

Back 12

∆H⊖f
The standard enthalpy of change of formation is the enthalpy change when one mole of a compound is formed from its elements under standard conditions.

Front 13

Define: the standard enthalphy of neutralisation.

Back 13

∆H⊖neut
The standard enthalpy of neutralisation is the enthalpy change when an acid and a base react to form one mole of water under standard conditions.

Front 14

Define: Bond enthalpy.

Back 14

Bond enthalpy (E) is the energy required to break one mole of a particluar type of bond in gaseous molecules under standard conditions. Standard conditions being 298K, 100kPa.

Front 15

Define: Hess's law.

(give a formula)

Back 15

Hess's law states the total enthalpy change for a reaction is the same for any reaction route, provided that the starting and finishing conditons are the same.

∆Hfθ = ∆H1 - ∆H2


∆H1 = ∆H2 + ∆H3

Front 16

Define: electrophile.

Back 16

An electrophile is an electron pair acceptor.

Front 17

Define: empirical formula.

Back 17

The emperical formula is the simplest whole-number ratio of the number of atoms of each element in a compound.

Front 18

Define: isotope.

Back 18

An isotope is an atom of the same element (with the same atomic number thus) that has a different mass. Isotopes have the same number of protons but different number of neutrons.

Front 19

Define: relative atomic mass.

Back 19

The relative atomic mass (Ar) of an element is the average relative mass of its atoms -taking into account its isotopes and their abundances - compared to one atom of C-12. Its the mass of 1 mole of atoms of an element.

Front 20

Define: cation.

Back 20

A cation is a positive ion formed from the loss of electon(s).

Front 21

Define: anion.

Back 21

An anion is a negative ion formed by the gain of electron(s).

Front 22

Define: ionic bond.

Back 22

An ionic bond is the electrostatic attraction between positive and negative ions.

Front 23

Define: Second electron affinity.

Back 23

The second electron affinity is the energy needed to be taken in (its endothermic thus) in order to add a second electron to an atom.

Front 24

Metal + acid -> ?

Back 24

metal + acid -> salt + hydrogen

example:
Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)

Front 25

Metal oxide + acid -> ?

Back 25

metal oxide + acid -> salt + water

example:
MgO(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l)

Why? Metal oxides are basic and bases neutralise acids ^^.

Front 26

metal carbonate + acid -> ?

Back 26

metal carbonate + acid -> salt + water + carbon dioxide.

example:
MgCO3(s) + 2HCl(aq) -> MgCl2(aq) + H2O(l) + CO2(g)

Front 27

metal hydroxide + acid -> ?

Back 27

metal hydroxide + acid -> salt + water.

example: of an important reaction:
NaOH(aq) + HCl(aq) -> NaCl(aq) + H2O(l)

Front 28

What is an alkali?

Back 28

Alkali's are soluble bases.

Front 29

What is the ionic equation for neutralisation?

Back 29

H+(aq) + OH-(aq) -> H2O(l)

Front 30

An acid does what when it reacts? Therefore an acid is...?

Back 30

Acids give up H+ ions when they react. Therefore acids are proton donors.

Front 31

Which of the following is a strong acid, and which is a weak acid?

*Hydrochloric acid
*Ethanoic acid

Back 31

Hydrochloric acid is a strong acid.

Ethanoic acid is a weak acid.

Front 32

What happens to a strong acid as it dissolves in water?

Back 32

A strong acid completely dissociates into ions when it dissolves in water.

Front 33

What happens to a weak acid as it dissolves in water?

Back 33

A weak acid only partially dissoiciates into ions when it dissolves in water.

Front 34

What is the difference between the strength of an acid and its concentration?

Back 34

The strength refers to the type of acid: and each type of acid can be diluted to different concentrations. Therefore a strong acid can be of lesser concentration than a weak acid, and yet it is still a strong acid.

(Note: However to compare the types of acids fairly it is best to use the same concentration.)

Front 35

If an acid reaches equilibrium does that mean it is a weak acid or a strong acid?

Back 35

A weak acid, as it is able to reach equilibrium because the acid does not dissociate into ions completely (so there is always leftover reactant).

Example:
ethanoic acid:

CH3COOH(aq) <=> H+(aq) + CH3COO-(aq)

Front 36

Compare a strong acid with a weak acid.

Back 36

* A strong acid completely dissociates into ions when dissolved in water (weak acids only partially).

*Strong acids have higher electrical conductivity than weak acids.

* Strong acids have a low pH (1 or 2) and weak have a higher pH (3-6).

* Strong acids are very corrosive (weak acids are less corrosive).

* Strong acids react more quickly than weak acids.

Front 37

If the electronic configuration of an element is 1s^2 what is meant by 1, s and 2?

Back 37

1 is the principal quantum number; s is the subshell; and 1 is the number of electrons in the subshell.

Front 38

What are three factors that affect ionisation energy?

Back 38

* atomic radius
* nuclear charge
* electron shielding

Front 39

What instrument is used to measure relative atomic mass?

Back 39

Mass spectrometer.

Front 40

How would you show that ammonia gas was given off?

Back 40

It turns damp litmus paper blue, and can also be recognized by its pungent smell.

Front 41

What is the volume of one mole of gas at r.t.p.?

Back 41

24 dm³

Front 42

State Le Chatelier's principle.

Back 42

If a change in conditions is made to a closed system in equilibrium, the system responds to counteract that change.

Front 43

state standard conditions.

Back 43

usually 100kPa and 298K.

Front 44

Name the structure and list the properties of a metal and state why it has these properties.

Back 44

Metallic bonding.

*High melting point -strong bonding due to delocalised electrons.
*can be worked into different shapes - layers of metal ions can move over eachother.
*high densitites - metal ions are packed close together.
*good electrical conductors - electrons are delocalised and can flow.

Front 45

Name the structure of NaCl and list the properties and state why it has these properties.

Back 45

Ionic structure / ionic bonding. -Giant 3D lattice.

*high melting poins & boiling points - due to strong electrostatic attraction between positive and negative ions.
*often soluble in water - ions interact with H2O dipoles.
*Conduct electricity in molten + aqueous state (not as solids!) - ions can move and carry charge in these states.

Front 46

What is the structure of I2 (Iodine), what are the properties of the structure, and state why it has these properties.

Back 46

Simple molecular structure - covalent bonding.

*very low boil/melting points - due to weak intermolecular forces.
*does not conduct electricity - bonding involves shared electrons (thus no charged particles)
*very low densities - they are usually gases or liquids due to weak I.M.F.s

Front 47

What is the structure of polymers or silicon dioxide? List the properties of this structure and state why these properties exist.

Back 47

Giant Molecular (Covalent) - covalent bonding.

*generally do not conduct electricity
*can be very hard - they have strong covalent bonds
*high melting/boiling points - due to these strong covalent bonds.
*

Front 48

What is the structure of graphite. List graphite's properties and state why it has these properties.

Back 48

Giant molecular (covalent) structure.

*allotrope of carbon
*conducts electricity - due to delocalised electrons between layers.
*covalent bonding within layers and thus strong there, but only weak i.m.f.s between layers (induced dipole-induced dipole) so layers break off easily.

Front 49

What structure does diamond have, list diamond's properties and state why it has these properties.

Back 49

Giant molecular (covalent) structure.

*Diamond is an allotrope of carbon.
*has very high melting/boiling points - due to very strong covalent bonding.
*insoluble in water - no charges to interact with polar water molecules.
*very hard - each atom is strongly bonded to 4 others.
*does not conduct electricity - complete covalent bonding and thus no charge movement.

Front 50

How does a mass spectrometer work?

Back 50

* Ionisation: the gaseous atoms are bombarded with electrons - which knock out an electron per atom - forming postively charged ions.

*Acceleration - accelerated by an electric field.

*Deflection - deflected by a magnetic field - lightest ions are the most deflected.

*Detection: the detector detects ions and their masses.

*info is printed out (graph).

Front 51

What is the maximum amount of electrons in a "s" subshell?

Back 51

2

Front 52

What is the maximum amount of electrons in a "p" subshell?

Back 52

6

Front 53

What is the maximum amount of electrons in a "d" subshell?

Back 53

10

Front 54

What is the maximum amount of electrons in a "f" subshell?

Back 54

14

Front 55

What is the order of repulsion with regard to electrons?

Back 55

1. lone pair-lone pair (greatest repulsion)
2. bonding pair-lone pair
3. bonding pair-bonding pair (least repulsion - but still quite a bit of repulsion)

Front 56

What is the most electronegative element?

Back 56

Flourine (4.0)

Front 57

concentration = ?

Back 57

concentration = (mass in grams) / (volume in dm^3)

or

concentration = (# of moles) / (volume in dm^3)

Front 58

What are i.m.f.s?

Back 58

Intermolecular forces are attractions (very weak bonding) between simple covalent molecules.

Front 59

List the three types of i.m.f.s?

Back 59

*induced dipole - induced dipole (aka Van der Waals)
*permanent dipole-permanent dipole
*hydrogen bonding (the strongest type of i.m.f.)

Front 60

Properties of polar molecules?

Back 60

*permanent dipole foces happen between polar molecules.
*higher boiling points

Happens due to large difference in electronegativity between elements in the molecule. - electrons get pulled toward the more electronegative element.

Front 61

What are properties of induced dipole forces

Back 61

*The bigger the atoms the more electrons (bigger electron cloud) thus the stronger the force (and more easily it can be polarised).
*They are the weakest type of i.m.f.

(Note: Induced ,temporary and instantaneous dipole all mean the same thing.)
*

Front 62

Why is ice less dense than liquid water?

Back 62

This is due to hydrogen bonding. Ice is less dencse than liquid water because the molecules are held further apart in a lattice of hydrogen bonds (some of these break on melting).

Front 63

Change of state: when do i.m.f.s break?

Back 63

*They partially break when in liquid state - molecules are allowed to move over eachother in liquid state.
*In gase state the i.m.f.s are broken completely as molecules seperate to form gas. (That's why boiling takes a lot of energy input.)

Front 64

List trends in electron configuration down a group and state why they exist.

Back 64

* number of protons and electrons increase - thus increase in atomic radius.
* more inner electron shells - thus more shielding of the nucleus -thus less attractive force between outer electrons and the nucleus so:
> first ionisation decreases
> electronegativity decreases

Front 65

List trends in electron configuration across a period and state why they exist.

Back 65

* Number of protons and electrons increase
* newest electrons all enter the smame shell: so there is an increase in positive charge but no increase in electron shielding as there are no new shells. (thus also not new inner shells)
*Atomic radius decreases
*First ionisation energy increases
*electronegativity increases (more nuclear charge - no extra shielding).

Front 66

Group 2 metals with oxygen: list key points.

Back 66

*tarnish in air
*burn vigorously when heated in air
*all form ionic metal oxides when they react with oxygen.

Front 67

Group 2 metals with water: state key points.

Back 67

*Be and Mg do not react (though Mg does with steam).
*Metals form hydrogen and a metal hydroxide.
*rate of reaction increases down the group. - as ionisation energy decreases.
*pH increases during reaction.

Front 68

What is used to test for CO2? Give equation.

What happens if too much CO2 is added?

Back 68

Lime water: calcium hydroxide.

Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)

(turns cloudy as CaCO3 is formed)

If too much CO2 is added then the the precipitate redissolves.

CaCO3(s) + CO2(g) + H2O(l) -> Ca(HCO3)2(aq)

Front 69

What are the colors of Chlorine, Bromine, and Iodine?

Back 69

*Chlorine = green
*Bromine = orange volatile liquid
*Iodine = grey/purple solid.

Front 70

How can you tell the difference between the halogens?

Back 70

Chlorine is a pale green solution. Bromine and iodine are both brown ,however, so:
Use a few drops of hexane.
*Bromine = orange/red layer
*Iodine = purple layer

Front 71

What is respiration?

Back 71

Respiration is another example of an important oxidation reaction that is exothermic. It is similar to combustion as CO2(g) and H2O(l) are the products. (And oxygen is one of the reactants.)

Front 72

What is the formuala for the enthalpy change of water?
What is the specific heat capacity of water?

Back 72

enthalpy change (q) = m x c x ∆T

m = mass
c = specific heat capacity of water which is 4.2 J/gK
∆T = change in temperature

Front 73

What is equilibrium?

And what do you see?

Back 73

Equilibrium is the state of balance where the concentrations (or pressures) of reactants and products remain constant. There is both a forward and backward reaction going on continuously.

At a microscopic scale the equilibrium is dynamic (constantly changing).


Macroscopically nothing can be seen as the forward and backward reaction are happening at the same rate so no overall change can be seen. Pressure/concentration stay constant.

Front 74

What is cracking and what are two different types of cracking?

Back 74

Cracking splits longer chain molecules into more useful shorter chain molecules.

*Thermal cracking - uses heat to split
*Catalytic cracking - uses a catalyst to split

Front 75

How can you tell there is a double bond present?

Back 75

Add bromine water (aqueous bromine). The solution will turn from orange to colorless if a double bond is present. (addition reaction)

Front 76

what is heterolytic fission?

Back 76

A breaking of a bond in which both electrons go to one atom. This makes a negatively charge ion.

Homolytic fission is when one electron goes to one atom, and the other electron goes to the other atom.

Front 77

What are nucleophiles?

Back 77

Nucleophiles molecules that donate a lone pair to a positively charged carbon atom. They are attracted to positive charge.

Front 78

How do you tell Chloride, bromide and iodide apart?

Back 78

By adding acidified AgNO2(aq). This forms a precipate. The faster the precipitate forms the faster the rate of reaction.

Chloride = white
bromide = cream
iodide = yellow

Front 79

What are/were CFCs used for and why?

Back 79

*Refrigerants
*Propellants
*Solvents
etc.

They seemed ideal due to their unreactive, non-flammable, boiling point being about room temperature and non-toxic (and dont cause corrosion) properties.

Front 80

Why are CFCs bad?

Back 80

They are responcible for the depletion of ozone in the stratosphere.

Because they are unreactive they reach the stratosphere rather than breaking up in the lower atmosphere. (They are so unreactive due to the C-F bond being extremely strong.) The UV radiation finally does break the bond and this produces Cl* free radicals. These catalyse the breakdown of the ozone.

Cl* + O3 → ClO* + O2
ClO* + O3 → Cl* + 2 O2
etc. Takes a very long time before the termination reaction takes place.

more saturated molecules are more reactive and thus somewhat better (used as replacements) such as HFCs and HCFCs.

Front 81

What is the formula for the ideal gas law?

Back 81

p x V = n x R x T

p = pressure (Pa)
V = volume (m^3)
n = # of gas moles
R = gas constant = 8.31 J/K mol
T = temperature (K)


assuming its an ideal gas.

Front 82

Shapes:

2 bonds (no lone pairs)

Back 82

Linear (thus planar)

Front 83

Shapes:

2 bonds (lone pair)

Back 83

V-shaped (non-planar)

Front 84

Shapes:

3 bonds (no lone pairs)

Back 84

Trigonal planar

Front 85

Shapes:

3 bonds (lone pair)

Back 85

pyrimidial (non planar)

Front 86

Shapes:

4 bonds (no lone pairs)

Back 86

Tetrahedral (non planar)

Front 87

Shapes:


5 bonds (no lone pairs)

Back 87

triagonal bipyramidal

Front 88

Shapes:

6 bonds (no lone pairs)

Back 88

octahedral

Front 89

reaction rate = ?

Back 89

reaction rate = change in amount or concentration / time taken

Front 90

rate equation:
rate = ?

Back 90

rate = K[X]^n

n = reaction order
K = constant
[X] = reactant

Front 91

Avogadro's constant = ?

Back 91

6 x 10^23

Front 92

Kc = ?

Back 92

Kc =

([C]^c) x ([D]^d)
--------------------
([A]^a) x ([B]^b)

A = compound A and a = moles of compound A

Front 93

Kp = ?

Back 93

Kp =


(pC^c) x (pD^d)
--------------------
(pA^a) x (pB^b)


p = partial pressure of that compound

A = compound A

a = moles of compound A

Front 94

What is the relative mass of a neutron?

Proton?

Electron?

Back 94

neutron: 1

proton: 1

electron: negligible

Front 95

What is the difference between a covalent bond and a dative covalent bond?

Back 95

The covalent bond is where two electrons are shared - one from each atom; and a dative covalent bond is when both shared electrons come from one atom.