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chem exam questions > A1 > Flashcards

A1 Flashcards

atomic structure, isotopes, electrons, ionic bonding

1
Q

isotope

A

atoms of the same element
same number of protons
different number of neutrons

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2
Q

element

A

substance made up of only 1 type of atom
cant be broken down into another substance

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3
Q

atom

A

smallest particle of an element that can exist

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4
Q

Relative masses of:
- proton
- neutron
- electron

A

1
1
1/1836

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5
Q

relative charges of:
- proton
- neutron
- electron

A

+1
0
-1

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6
Q

atomic number

A

the number of protons

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7
Q

mass number

A

sum of protons + neutrons

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8
Q

why do isotopes have the same chemical properties

A

they have the same number of electrons

same arrangement of electrons
neutrons have no effect on chemical properties

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9
Q

anion

A

negative ion
formed when atom GAINS electrons

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10
Q

cation

A

positive ion
formed when atom LOSES electrons

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11
Q

relative atomic mass

A

weighted mean mass
of an atom of an element
compared to 1/12th of the mass of an atom of carbon 12

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12
Q

relative isotopic mass

A

weighted mean mass
of an atom of an isotope
compared to 1/12th of the mass of an atom of carbon 12

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13
Q

binary compound

A

contains only 2 elements

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14
Q

polyatomic ions

A

ions which have more than one type of atom

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15
Q

how many electrons do these shells hold?
- 1st shell
- 2nd shell
- 3rd shell
- 4th shell

A

2
8
18
32

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16
Q

how many electrons do these subshells hold?
- s
- p
- d
- f

A

2
6
10
14

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17
Q

why do we write 4s before 3d

A

because 4s is in a lower energy level than 3d
we write starting with lowest energy level

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18
Q

orbital

A

region within an atom
around the nucleus
which holds up to 2 electrons with OPPOSITE SPINS

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19
Q

what do the arrows represent in electrons in boxes

A

spin

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20
Q

give the shape of and s orbital and a p orbital

A

spherical
dumbell

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21
Q

what are the 2 exceptions to electron configurations

A

chromium
copper

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21
Q

ionic bonding

A

electrostatic attraction
between oppositely charged ions

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21
Q

describe structure and physical properties of ionic bonding

A
  • giant ionic lattice
  • high melting/boiling points
  • soluble in POLAR solvents
  • solid at room temperature
  • brittle
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22
Q

what 2 factors affect ionic bonding strength

A

ionic charges
ionic radii

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23
explain how ionic charges and radii size affect ionic bonding
stronger ionic charges= stronger attraction= stronger ionic bond smaller radii= stronger attraction= stronger ionic bond
24
how does an ionic compound dissolve in water
ionic lattice breaks down the ions separate water molecules surround ions
25
why don't ionic compounds conduct electricity when solid
ions are in fixed position and cannot move through the structure
26
why do ionic compounds conduct electricity when molten/ aqueous
ions are mobile are free to move through structure and carry charge
27
define molten
liquified by heat
28
define aqueous
dissolved in water
29
why do ionic compounds have high boiling/melting points
- giant ionic lattice - strong ionic bonds - these break during melting - require lots of energy to overcome
30
covalent bonding
shared pair of electrons electrostatic attraction between shared pair of electrons and positive nuclei
31
covalent bond is localised- what does this mean?
the attraction is soley between the shared pair of electrons and nuclei of the bonded atoms
32
difference between ionic and covalent bonding
in ionic, electrons are transferred, in covalent electrons are shared
33
lone pair
a pair of electrons that is not shared
34
multiple bond
2 atoms share more than 1 pair of electrons
34
double bond
attraction is between 2 shared pairs of electrons and nuclei of bonding atoms
35
triple bond
attraction is between 3 shared pairs of electrons and nuclei of bonding atoms
36
dative bond
both electrons in the shared pair come from 1 of the bonding atoms
37
properties of simple molecular substances
- low melting/boiling points (weak IMFs require little energy to overcome) -dissolve in NON-POLAR solvents (IMFs form between non-polar solvent and simple molecule) -don't conduct electricity (no free electrons/ ions) -gases
38
properties of giant covalent structures
- high melting/boiling points (strong covalent bonds require high energy to overcome) -solids -insoluble in polar and non-polar solvents (covalent bonds too strong) - don't conduct electricity (except graphite)
39
why does graphite conduct electricity
only uses 3 of its 4 bonds has a free, delocalised electron to move through the structure and carry charge
40
electron pair repulsion theory
electrons have negative charges so electron pairs repel each other as far as possible lone pairs repel more than bonding pairs as they're more electron dense (decrease bond angle by 2.5)
41
octet rule
not all elements achieve a full outer shell of 8 - may be too few/much electrons elements beyond group 15 and period 3 can expand their octet
42
why are non linear shapes and trigonal bipyramids POLAR
-dipoles are not symmetrically arranged - effects don't cancel out -there is an overall dipole
43
state bond angles within -2 regions -3 regions -4 regions
2 regions= 180 degrees 3 regions= 120 degrees 4 regions= 109.5 degrees
44
electronegativity
the ability of a bonded atom to attract the shared pair of electrons in a covalent bond
45
how is electronegativity measured
pauling scale
46
how does number of protons and atomic radius affect electronegativity
more protons= increase electronegativity (electrostatic attraction towards electrons is greater) small distance between outer shell and nucleus= increase electronegativity (stronger attraction)
47
why are noble gases not electronegative
they have a full outer shell of electrons don't attract electrons
48
what makes a bond non-polar? (pure covalent bond)
both atoms have same electronegativities electron pair is shared evenly
49
what makes a bond polar?
one atom is more electronegative it attracts the shared pair of electrons the shared pair is unevenly shared
50
why are C--H bonds non-polar?
carbon and hydrogen have very similiar electronegativites
51
in MgI there is polarisation- explain this term (3)
- Mg and I have different electronegativities - electron pair is shared unevenly - there is an uneven distribution of charge
52
state the 3 types of intermolecular forces
-induced dipole-dipole -permanent dipole-dipole -hydrogen bonding
53
describe the origin of induced dipoles (london forces)
electrons in the electron cloud are constantly moving at any moment, there is an instantaneous dipole more electrons on one side this instantaneous dipole induces a temporary dipole in a neighbouring molecule
54
describe how number of electrons affects strength of london forces
more electrons means a larger instantaneous dipole and larger induced dipole stronger attractive forces between molecules
55
what does a hydrogen bond act between
lone pair on an electronegative atom and a hydrogen atom
56
state anomolous properties of water
ice is less dense than water so it floats relatively high Mp/Bp high surface tension
57
explain why ice floats on water
ice is less dense than water because hydrogen bonds hold the water molecules apart in an open lattice each water molecule can form 4 hydrogen bonds as there are 2 lone pairs on the O and 2 H atoms
58
explain why water has a relatively high melting/boiling point
relatively strong hydrogen bonds which require large amount of energy to overcome
59
why does boiling point increase going down the noble gases
there are more electrons so stronger london forces more energy required to overcome
60
state the formula and charge of: -nitrite -hydrogencarbonate -manganate(VI) -silver
NO2- HCO3- MnO4- Ag+
61
state the formula and charge of: -sulfite -dichromate -phosphate -sulfide
SO3 2- Cr2O7 2- PO4 3- S 2-
62
metallic bonding
electrostatic attraction between cations and delocalised electrons
63
describe the properties of metals
high Mp/Bp (strong metallic bonds) conduct electricity (delocalised electrons) insoluble (strong metallic bonds) solid at room temp malleable (layer of cations slide over each other) ductile giant metallic lattice

chem exam questions (11 decks)

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