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A LEVEL CHEM > MOD. 2 chap 2.1-2.2 > Flashcards

MOD. 2 chap 2.1-2.2 Flashcards

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1
Q

define relative atomic mass

A

the weighted mean mass of an atom of an element compared with one-twelfth of the mass of an atom of carbon-12

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2
Q

define relative isotopic mass

A

the mass of an atom of an isotope compared with one-twelfth of the mass of an atom of carbon-12

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3
Q

How to calculate relative atomic mass

A

mass of isotope*abundance
divide by 100
round

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4
Q

Name scientists in order of their discoveries regarding the atom

A

Democritus
Dalton
J.J. Thomson-discovers electrons
Rutherford-gold leaf experiment
Niels Bohr
Rutherford discovers proton
Chadwick discovers neutron

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5
Q

define atomic number

A

number of protons

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6
Q

define mass number

A

number of neutrons and protons

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7
Q

atom economy formula

A

molecular mass of desired product/molecular mass of products *100

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8
Q

percentage yield formula

A

actual amount in mol of products/theoretical amount in mol of product *100

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9
Q

what does stoichiometry tell us

A

the amounts of substances that are involved in chemical reactions

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10
Q

name acids and their formulas

A

Sulfuric acid H2SO4
Hydrochloric acid HCl
Nitric acid HNO3

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11
Q

Name Alkalis and their formulas

A

sodium hydroxide NaOH
Potassium Hydroxide KOH
Ammonia NH3

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12
Q

Name bases and their formulas

A

Metal Oxides
Metal Hydroxides
Ammonia(NH3)
Amines

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13
Q

what is a proton acceptor

A

A base

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14
Q

what is proton donor

A

an acid they release H+ ions into the solution

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15
Q

State formula for ammonia reaction with water to form a weak base

A

NH3 (aq)+H2O(l)–> reversible NH4+(aq)+OH-

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16
Q

why is ammonia a weak base

A

only a small proportion of the dissolved NH3 reacts with water. that is also why it is a reversible reaction

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17
Q

what are amphoteric substances+ example

A

substances that can behave like both acids and bases.
for example amino acid molecules:

carboxyl acid group COOH can donate a proton and the amino basic group can accept the proton.

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18
Q

4 ways to make a salt

A
  1. acid+ metal carbonates–> salt+H2O+CO2
  2. acid+ metal oxide–> salt+H2O
  3. acid+ alkali–> salt+ H2O
  4. acid+ metal–> salt+ hydrogen
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19
Q

what is formed when acids are neutralised by aqueous ammonia
ammonium nitrite+ Nitric acid–>???

A

ammonium salts are formed

NH3(aq) +HNO3(aq) —> NH4NO3(aq)

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20
Q

Ionic equation for neutralisation

A

H+(aq) +OH-(aq) —-> H2O(l)

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21
Q

define hydrated compound

A

crystalline form containing water

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22
Q

define anhydrous compound

A

a form containing no water

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23
Q

define water of crystallisation

A

water molecules that form an essential part of the crystalline structure of a compound.

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24
Q

Methyl Orange colour in acid base and the end colour

A

acid- red
base-orange
End colour- orange

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25
Q

Phenolphthalein colour in acid base and the end colour

A

acid-colourless
base-pink
end colour-pale pink

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26
Q

bromothymol blue colour in acid base and the end colour

A

acid- yellow
base-blue
end colour- green

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27
Q

How many electrons in shell
1
2
3
4

A

1st shell=2 e-
2nd shell=8 e-
3rd shell=18 e-
4th shell=32 e-

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28
Q

define the term shell

A

a group of atomic orbitals with the same principal quantum number, n.

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29
Q

describe what is meant by principal quantum number

A

the number representing the overall energy of each orbital, which increases with distance from the nucleus. sets of orbitals that have the same n value are referred to as electron shells/energy shells

30
Q

state the 4 different types of orbitals

A

s p d f

31
Q

what is the principal quantum number of shell 2

A

n=2

32
Q

Max number of electrons formula

A

2n^2

33
Q

state what is meant by an atomic orbital

A

the region around the nucleus that can hold up to 2 electrons with opposite spin
electrons can have either an up spin or down spin.

34
Q

what is an electron NOT GCSE

A

not a particle, it is considered a cloud of negative charge

35
Q

ionic bonding happens between

A

a metal and a non metal.
electrons are transferred from the metal to the non metal to form oppositely charged ions which attract one another, they are bonded together by electrostatic forces of attraction.

the metal ion is positive
the non metal is negative

36
Q

Describe covalent bonding

A

occurs between compounds consisting of 2 non metals.
Negatively charged pair of electrons is attracted to the positive charges of both nuclei. This attraction overcomes the repulsion between the 2 positively charged nuclei. Electron pair occupies space between the 2 atoms nuclei. So covalent bond holds the 2 atoms together.

37
Q

what is metallic bonding

A

occurs in metals. Electrons are shared between all the atoms

38
Q

Describe the structure of giant ionic lattices

A

each ion is surrounded by oppositely charged ions, these ions attract each other from all directions forming giant lattices with huge amounts pf electrostatic forces of attraction between ions

39
Q

describe structure and bonding of sodium chloride

A

each Na+ ion is surrounded by 6 Cl- ions
each Cl- ion is surrounded by 6Na+ ions

40
Q

State and explain properties of ionic compounds

A

*HIGH MELTING AND BOILING POINT
-ionic compounds at room temp.= SOLID
-lots of energy needed to break the strong electrostatic bonds that hold oppositely charged ions together.
—> the greater the charge the stronger the electrostatic forces of attraction

*ELECTRICAL CONDUCTIVITY
-when solid ions are held in fixed positions and they cannot move, therefore the ionic compound will not conduct electricity
BUT if molten/dissolved in water, ions are free to move and carry charge (can conduct electricity).

*SOLUBILITY
-ionic lattice dissolves in polar solvents (Polar bonds occur between atoms that do not share electrons equally therefore atoms have very small charges on them). These polar solvents disturb ionic lattice by attracting charged ions in the giant lattice pulling them out.

41
Q

Describe what is meant by single covalent bonding

A

Only one shared pair of electrons in a single bond acting in only one direction only between the atoms involved in the bond.

42
Q

How many covalent bonds do the following make:
carbon
nitrogen
oxygen
hydrogen

A

carbon- 4 covalent bonds
nitrogen- 3 covalent bonds
oxygen- 2 covalent bonds
hydrogen-1 covalent bond

43
Q

What do we mean by lone pairs

A

pairs of electrons in the outer shell which are not involved in chemical bonding

44
Q

what do we call the average energy needed to break bonds

A

average bond enthalpy

45
Q

define dative covalent bonding

A

where one of the atoms supplies both the shared electrons to the covalent bond

46
Q

Identify the 2 types of covalent bonding

A

simple molecular lattice
giant covalent lattice

47
Q

Describe what is meant by simple molecular lattices

A

made of simple molecules
atoms within each molecule are held by strong covalent bonds. Different molecules are held by weak intermolecular forces (van der waals/london forces)

48
Q

Give an example of simple molecular lattice

A

Iodine

49
Q

describe the properties of simple molecular structures

A

*LOW MELTING AND BOILING POINTS
-intermolecular forces are weak, small amount of energy needed to break them

*NON CONDUCTORS OF ELECTRICITY
-no charged particles are free to move and carry charge

*SOLUBLE IN NON POLAR SOLVENTS
-Weak London forces can form between covalent molecules and solvents, this helps molecular lattice be broken down

50
Q

state examples of giant covalent structures

A

diamond, graphite, SiO2

51
Q

what determines the shape of compounds

A

number of electron pairs/densities around the central atom

52
Q

name of shape when 1 pair of bonded electron pairs are around the central atom

A

linear

53
Q

name of shape when 2 pairs of bonded electron pairs are around the central atom +angle

A

linear
angle=180

54
Q

name of shape when 3 pairs of bonded electron pairs are around the central atom

A

trigonal planar
angle=120

55
Q

name of shape when 4 pairs of bonded electron pairs are around the central atom + angle

A

tetrahedral
angle=109.5

56
Q

name of shape when 4 pairs of bonded electron pairs are around the central atom +angle

A

trigonal bipyramid
angle=120

57
Q

name of shape when 6 pairs of bonded electron pairs are around the central atom +angle

A

octahedral
angle=90

58
Q

which one is more electron dense- lone pairs or bonded pairs

A

lone pairs are more electron dense

59
Q

define electronegativity

A

measures the attraction of a bonded atom for the pair of electrons in a covalent bond
-the ability of an atom to attract the pair of electrons in a covalent bond

60
Q

name the factors affecting electronegativity

A

the number of protons in the nucleus; having more protons increases the electronegativity

the distance from the nucleus;

the amount of screening by inner electrons.

60
Q

what happens to electronegativity going down the group

A

As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.

60
Q

why is ice less dense than water

A

when ice forms water molecules arrange themselves into an orderly pattern and hydrogen bonds form between molecules

Ice has an open lattice with hydrogen bonds holding water molecules apart.

60
Q

what are London /dispersion forces caused by

A

constant random movement of electrons in atom’s shells. This movement unbalances the distribution of charge within the electron shells

61
Q

why is does water have a meting point higher than expected

A

hydrogen bonds are much longer than other intermolecular forces.

lots of energy required to overcome these strong forces of attraction

61
Q

Key Concepts of VSEPR- vallance shell electron pair repulsion theory

A

Electron pairs around a central atom (including both bonding pairs and lone pairs) repel each other.

The shape of a molecule is determined by the arrangement of these electron pairs to minimize repulsion.

Bonding pairs of electrons form bonds between atoms, while lone pairs (non-bonding pairs) are unshared pairs of electrons on the central atom. Both types affect the molecular shape.

Lone pairs exert more repulsion than bonding pairs, which can lead to deviations in ideal bond angles.

62
Q

rules for hydrogen bonding

A

must bond to a very electronegative atom - F O N that has a lone pair of electrons

a hydrogen atom covalently bonded to this electronegative atom

63
Q

define oxidation numbers

A

the number of electrons lost or gained or shared by an atoms in a compound.

64
Q

what is the oxidation number for uncombined elements

A

oxidation number 0 because they have not gained or shared or lost any electrons

65
Q

difference between inter and intramolecular forces

A

Intermolecular forces are the electrostatic interactions between molecules. The intermolecular forces are usually much weaker than the intramolecular forces

Intramolecular forces are the chemical bonds holding the atoms together in the molecules. The three major types of chemical bonds are the metallic bond, the ionic bond, and the covalent bond.

66
Q

When acid is added to water, water molecules form _____ ions

A

Oxonium ions H3O+

A LEVEL CHEM (10 decks)

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