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

MOD. 2 chap 2.1-2.2 Flashcards

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
Phenolphthalein colour in acid base and the end colour
acid-colourless base-pink end colour-pale pink
26
bromothymol blue colour in acid base and the end colour
acid- yellow base-blue end colour- green
27
How many electrons in shell 1 2 3 4
1st shell=2 e- 2nd shell=8 e- 3rd shell=18 e- 4th shell=32 e-
28
define the term shell
a group of atomic orbitals with the same principal quantum number, n.
29
describe what is meant by principal quantum number
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
state the 4 different types of orbitals
s p d f
31
what is the principal quantum number of shell 2
n=2
32
Max number of electrons formula
2n^2
33
state what is meant by an atomic orbital
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
what is an electron NOT GCSE
not a particle, it is considered a cloud of negative charge
35
ionic bonding happens between
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
Describe covalent bonding
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
what is metallic bonding
occurs in metals. Electrons are shared between all the atoms
38
Describe the structure of giant ionic lattices
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
describe structure and bonding of sodium chloride
each Na+ ion is surrounded by 6 Cl- ions each Cl- ion is surrounded by 6Na+ ions
40
State and explain properties of ionic compounds
*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
Describe what is meant by single covalent bonding
Only one shared pair of electrons in a single bond acting in only one direction only between the atoms involved in the bond.
42
How many covalent bonds do the following make: carbon nitrogen oxygen hydrogen
carbon- 4 covalent bonds nitrogen- 3 covalent bonds oxygen- 2 covalent bonds hydrogen-1 covalent bond
43
What do we mean by lone pairs
pairs of electrons in the outer shell which are not involved in chemical bonding
44
what do we call the average energy needed to break bonds
average bond enthalpy
45
define dative covalent bonding
where one of the atoms supplies both the shared electrons to the covalent bond
46
Identify the 2 types of covalent bonding
simple molecular lattice giant covalent lattice
47
Describe what is meant by simple molecular lattices
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
Give an example of simple molecular lattice
Iodine
49
describe the properties of simple molecular structures
*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
state examples of giant covalent structures
diamond, graphite, SiO2
51
what determines the shape of compounds
number of electron pairs/densities around the central atom
52
name of shape when 1 pair of bonded electron pairs are around the central atom
linear
53
name of shape when 2 pairs of bonded electron pairs are around the central atom +angle
linear angle=180
54
name of shape when 3 pairs of bonded electron pairs are around the central atom
trigonal planar angle=120
55
name of shape when 4 pairs of bonded electron pairs are around the central atom + angle
tetrahedral angle=109.5
56
name of shape when 4 pairs of bonded electron pairs are around the central atom +angle
trigonal bipyramid angle=120
57
name of shape when 6 pairs of bonded electron pairs are around the central atom +angle
octahedral angle=90
58
which one is more electron dense- lone pairs or bonded pairs
lone pairs are more electron dense
59
define electronegativity
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
name the factors affecting electronegativity
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
what happens to electronegativity going down the group
As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.
60
why is ice less dense than water
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
what are London /dispersion forces caused by
constant random movement of electrons in atom's shells. This movement unbalances the distribution of charge within the electron shells
61
why is does water have a meting point higher than expected
hydrogen bonds are much longer than other intermolecular forces. lots of energy required to overcome these strong forces of attraction
61
Key Concepts of VSEPR- vallance shell electron pair repulsion theory
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
rules for hydrogen bonding
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
define oxidation numbers
the number of electrons lost or gained or shared by an atoms in a compound.
64
what is the oxidation number for uncombined elements
oxidation number 0 because they have not gained or shared or lost any electrons
65
difference between inter and intramolecular forces
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
When acid is added to water, water molecules form _____ ions
Oxonium ions H3O+

A LEVEL CHEM (12 decks)

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