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Prelim > ASB (atomic structure and bonding) > Flashcards

ASB (atomic structure and bonding) Flashcards

1
Q

define matter

A

Anything with mass that takes up volume(solid liquid gases etc)

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2
Q

relative masses, charges and locations of P, N and E

A

Relative masses: p = 1, n = 1, e = 0
Relative charges: p = +1, n = 0, e = -1
Location:
p in nucleus
n in nucleus
e orbiting in shells

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3
Q

what is atomic number and mass number

A

atomic number (z)= number of protons

mass number (A)= proton number + nuetron

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4
Q

define relative atomic mass (Ar)

A

the number of times heavier an average atom is than one-twelfth of a carbon-12 atom - taking into account the different isotopic masses and their relative abundance

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5
Q

equation for Ar

A

Relative Atomic Mass = (mass x % abundance) + (mass 2 x % abundance) / 100

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6
Q

define ion

A

a positively or negatively charged atom, or covalently bonded group of atoms

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7
Q

define isotopes - how are they different/same

A

Atoms with the same atomic number but different mass number/ the same number of protons but a different number of neutrons”
- reactivity is the same but physical properties are slightly different

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8
Q

define ionic bonding

A

occurs between a metal and non-metal atom, atoms lose or gain electrons to complete a shell

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9
Q

define covalent bonding

A

occurs between two non-metal atoms, electrons are shared to complete a full shell

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10
Q

define metallic bond

A

electrostatic attraction that exists between the positively charged cations and the sea of delocalised electrons, occurs between metals,

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11
Q

describe cations and anions

A

Cations- positive (cats have paws), losing electrons, metal ions

Anions- negative, gaining electrons, non metal ions

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12
Q

name molecular ions and their symbols

A

Ammonium- NH₄⁺

Hydroxide OH¯

carbonate CO₃²⁻

phosphate PO₄³⁻

Nitrate NO₃-

Sulfate SO₄²-

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13
Q

factors that affect force of attraction:

A
  • strength of bond (smaller = stronger)
  • charge of the iron (higher = stronger)

^ Together they’re called charge density

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14
Q

what are the 3 rules of octet and the exceptions

A

Rule of Octet

  1. Stability- Atoms are most stable when their outermost energy level is full
  2. ionic Bonding: Atoms may gain or lose electrons to achieve a full octet.
  3. Covalent Bonding:Atoms may also share electrons through covalent bonding to achieve a full octet.

Exceptions:

While the rule of octet applies to many molecules and ions, there are exceptions. Some atoms, particularly those in the second period (e.g., boron, beryllium, and hydrogen), may form stable molecules with fewer than eight electrons in their outer shell

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15
Q

Dative Bond

A

when both electrons are provided by the same atom (lone pair)

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16
Q

why do atoms form compounds

A

to become stable and gain a full outer shell (noble gas formation)

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17
Q

define valecy

A
  • number of other atoms that an atom of an element can combine with depends on the number of electrons it can lose, gain or share
  • i.e. the number of electrons in its outer shell
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18
Q

which elements dont normally form ionic compounds becuase of the large amount of energy required

A

C, Si, Be and B

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19
Q

complete the acid reactions-

  1. Acid + Alkali >
  2. Acid + Base >
  3. Acid + Carbonate >
  4. Acid + Metal >
A
  1. Salt + Water
  2. Salt + Water
  3. Salt + Water + Carbon dioxide
  4. Salt + Hydrogen
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20
Q

what is avogadros constant (mole)

A

6.022 × 10²³

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21
Q

what is the difference between Relative Molecular mass (Mr) and Relative Formula Mass (RFM)

A

Mr, of a compound is relative to C-12 and RFM is the AVERAGE mass of the formula of an ionic compound relative to C-12

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22
Q

How do you work out Mr and RFM

A

Add up all the Ar’s present,
e.g. RFM of NaOH = 23+ 16 + 1 = 40

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23
Q

formula to work out number of moles

A

Number of moles = mass (in grams) / Molar Mass (Ar, Mr, RFM)

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24
Q

Difference between Empirical Formula and Molecular formula

A

EF is the simplest whole-number ratio of elements present in a compound while MF is the actual whole-number ratio of elements present in a compound

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25
Q

how to find out empirical formula

A
  1. List all of the elements in the compound
  2. Write the experimental masses (if using mass data) or % (if using % composition) underneath them
  3. Divide each mass or % by the Ar of that element (to find the number of moles of that mass)
  4. Divide all of the numbers by the smallest (to get the molar ratio) Doing this gets the smallest number to 1 and all other numbers are relative to this
  5. Find the simplest whole number ratio
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26
Q

whats equivalent to 1dm³

A

1 litre, 1000cm³, 1000ml

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27
Q

what is conservation of matter

A
  • in a closed system, the total amount of matter (or mass) remains constant.
  • matter cannot be created or destroyed in ordinary chemical or physical processes; it can only change from one form to another.
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28
Q

what values are equal to 1dm3

A

1 dm3 ≡ 1 Litre≡ 1000 cm3 = 0.001m3

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29
Q

whats the properties of gas

A
  • Fill all of the space available to them
  • Expand upon heating
  • Exert pressure on the walls of their container
  • Pressure changes as temperature changes - more kinetic energy so more frequent collisions
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30
Q

ideal gas law equation

A

PV=NRT

  • P = pressure (Pa / kPa)
  • V = volume (m3 / dm3)
  • n = number of moles of gas
  • R = gas constant (8.314 JK-1 mol-1)
  • T = temperature (K) (add 273 to celcius)
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31
Q

differences between ideal gas and real gas

A

ideal

  • no intermolecular forces
  • molecules dont take up space- but have mass
  • temperature of the gas is related to the average kinetic energy of the molecules
  • Collisions between molecules are elastic (no loss of kinetic energy)

real

  • intermolecular forces- reduces pressure from ideal gas
  • molecules take up space- increases volume of ideal gas
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32
Q

gas volume equations (at room temp)

A

vol(cm3) = n x 24000

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33
Q

concentration equation

A

conc (mol dm-3) = mol/ vol (dm-3)

34
Q

Qualitative analytical chemistry

A
  • Determining the nature of the constituents of a sample.
  • eg. “a blue-green precipitate”, flame tests
35
Q

Quantitative analytical chemistry-

A
  • Determining how much of one or more constituents is present in a particular sample.
  • Numerical observation.
  • E.g. “Cholesterol levels were elevated by 0.5mmol dm-3 to 4.3mmol dm-3
36
Q

Gravimetric analysis-

A
  • Constituents being determined can be isolated in some weighable form.
  • Uses the force of Gravity (metric) Gravi
  • E.g. Filtering a pure precipitate and weighing
37
Q

Volumetric analysis-

A
  • The method used to determine the amount of a constituent involves measuring the volume of a reagent.
  • Uses volume (metric) volu
  • E.g. Titration
38
Q

what is a titration

A
  • 2 solutions mixed to reach the end point.
  • often is a neutralisation
39
Q

what are 4 types of titration

A
  • Acid-base
  • Redox titration
  • Complexometric
  • Back titration
40
Q

define ionisation energy and give the equation

A
  • the energy required to remove one mole of an electron from a gaseous atom.
41
Q

what are 3 factors affecting ionisation energy and describe them

A
  • nuclear charge- atoms with more protons= stronger positive charge= stronger electrostatic attraction between nucleus and electron
  • atomic radius- as distance increases between the nucleus and electrons the electrostatic attraction decreases
  • electron shielding- Inner electron shells shield the outermost electrons from the full attractive force of the nucleus so easier to be lost.
42
Q

what are the 4 sublevels and how many electrons can each hold

A

s- 2
p- 6
d- 10
f- 14

43
Q

what are the electronic configurations for K, Ca, Li, Cl, Mg

A

K- 1s2 2s2 2p6 3s2 3p6 4s1

Ca - 1s2 2s2 2p6 3s2 3p6 4s2

Li- 1s2 2s1

Cl- 1s2 2s2 2p6 3s2 3p5

Mg- 1s2 2s2 2p6 3s2 OR * Ne [3s2]

44
Q

what are the exceptions for electronic configuration

A

chromium- 1s2 2s2 2p6 3s2 3p6 4s1 3d5

copper- 1s2 2s2 2p6 3s2 3p6 4s1 3d10

45
Q

what is an orbital

A

Each energy sub-level has one or more orbitals, each of which can contain a maximum of two electrons

45
Q

describe the shapes of s and p sub levels

A

S- The s orbital has a very simple and symmetric shape, which can be described as a sphere.

P- distinctive figure-eight shape found either side of the nucleus

46
Q

how many orbits does each sublevel have

A

s= 2 electrons= 1 orbit

p= 6 electrons= 3 orbits

d- 10 electrons= 5 orbits

47
Q

define radioactive decay

A

spontaneous breakdown of a nucleus to form a more stable one

48
Q

give an example of alpha decay

A

238 234 4
U → Th + a
92 90 2

49
Q

give an example of beta decay

A

39 39. 0
K. →. Ca. +. e
19 20. -1

50
Q

give an example of gamma decay

A

99m 99
Tc. → Tc. + Y
43 43

m = metastable

51
Q

define half life

A

the time it takes for half the number of radioactive particles in a sample to decay

52
Q

describe T1 and T2 images

A

T1 image = fat is white

T2 image = both fat and water is white

53
Q

what is the electron pair repulsion theory

A

electron pairs are clouds of negative charge, so there is mutual repulsion between them, forcing them as far apart as possible.

54
Q

what is the shape of a molecule with 2 bonds (or with 1 lone pair)

A

Linear (non linear)

55
Q

what is the shape of a molecule with 3 bonds (or with 1 lone pair)

A

trigonal planar(1 lone - trigonal pyramidal)

56
Q

what is the shape of a molecule with 4 bonds (or with 1/2 lone pair)

A

tetrahedral (1 lone- pyramidal, 2 lone- v shaped)

57
Q

what is the shape of a molecule with 5 bonds (or with 1/2 lone pair)

A

trigonal bipyramid (1 lone- seesaw, 2 lone- t shaped)

58
Q

what is the shape of a molecule with 6 bonds (or with 1/2 lone pair)

A

octahedral (1 lone - square pyramid, 2 lone- square planar)

59
Q

why do lone pairs have more repulse between that than bonded pairs

A
  • Greater negative charge concentration
  • lone pairs also do not share their electrons between 2 nuclei
  • greater electron density
60
Q

define electronegativity

A

The ability of an atom to attract a bonding pair of electrons in a covalent bond’

61
Q

which elements are most and least electronegative and what is the pattern along periods and groups

A
  • electronegativity is measure in the pauling scale.
  • fluorine is most electronegative - 4.0
  • along a period it becomes more electronegative and down a group it becomes less electronegative
62
Q

what are the rules of electronegativity

A
  • Difference in electronegativity <0.5 = non-polar covalent
  • Difference in electronegativity between 0.5 and 1.6 = polar covalent
  • Difference in electronegativity >2.0 = ionic

If the difference in electronegativity is between 1.6 and 2.0…

  • If a metal is involved = ionic
  • If only non-metals are involved = polar covalent
63
Q

what is a polar bond + example

A

when atoms have slightly varying electronegativities, one may have a different charge to the other- a permenant dipole- e.g H-Cl

64
Q

what affects electronegativity

A
  • Number of protons in nucleus
  • Distance of outer electrons from nucleus
  • Amount of shielding by inner electrons
65
Q

compare intra and inter molecular forces

A

Intramolecular
- strong
- between atoms
-ionic/covalent/dative/ metallic

Intermolecular
- weak
- between molecules
- IDID/van der waals
- permanent dipole
- hydrogen

66
Q

Describe IDID forces

A
  • in molecules that are typically non-polar like H2, electrons are moving within their orbitals at very high speeds so more electrons may find themselves towards one end of the molecule. So temporarily that end will be δ- and the other end, δ+
  • if the instantaneous dipole is close to another non-polar molecule, the electrons may be repelled by the δ- to either, creating an induced dipole
  • they now have a Van der Waals attraction
  • easily broken
  • IDID force increase as the number of electrons increase down the group higher boiling point)
67
Q

Describe hydrogen bonds

A
  • strongest attraction
  • hydrogen atom is covalently bonded to one of the three most electronegative atoms (N, O, F)
  • NH3, H2O and HF have higher boiling points than expected.
67
Q

Describe Permenant dipole- dipole forces

A
  • attraction between molecules that already have a permanent dipole (where one atom has a higher electronegativity than the other)
  • harder to break
  • IDID still has an effect
68
Q

what kind of melting point do ionic binds have and why?

A

high melting point due to strong electrostatic attractions between the positive and negative ions

69
Q

strength of ionic bond depends on…

A
  • charges of the ion (Calcium Oxide with Ca2+ and O2- has a higher melting point than Potassium Chloride with K+ and Cl-)
  • size of the ion (smaller ions can pack more closely than larger ions & have a greater charge density on surface of ions, so the attractions are greater)
    size is greater factor
70
Q

electrical conductivity of ionic bond

A
  • solid ionic lattices don’t conduct electricity because they are fixed in a crystal lattice
  • when molten or dissolved the ions are free to move so can carry charge so can conduct electricity
71
Q

How do you draw metallic bonding

A

Metallic bond formed by electrostatic attraction between the cations and sea of electrons

72
Q

conductivity of metals

A
  • All metals conduct.
  • Delocalised electrons are free to move throughout the structure in 3-dimensions
73
Q

structure of metals

A
  • Metals are malleable (can be beaten into sheets) and ductile (can be pulled into wires).
  • Atoms can roll past each other without breaking the metallic bonds.
74
Q

melting point of giant molecular covalent bond

A

high boiling point because the covalent (intramolecular) bonds have to be broken

75
Q

structure of diamond and graphite (giant molevular covalent)

A

diamond:

  • Strong bonds hold lattice in rigid structure
  • Insoluble in water and organic solvents because they have no charge

graphite:

  • Common form of carbon found in soot, charcoal, pencil lead
  • Carbon atom joined to 3 other carbon atoms by strong covalent bonds. This creates hexagonal sheets of atoms
76
Q

melting point of simple covalent

A
  • when heated, the intermolecular hydrogen bonds are broken not the covalent bond
  • low melting point becuase the intermolecular forces arr weak
77
Q

describe group 2 elements

A
  • Alkaline earth metal oxides and hydroxides are basic
  • very reactive
  • all ionic compounds (beryllium has covalent character)
  • higher melting points and stronger metallic bonds than group 1
  • because they form 2+ ions rather than 1+ ions, releasing 2 delocalised electrons rather than 1. 2 electrons will carry more charge than 1 electron per ion
78
Q

how does atomic number impact IDID forces

A

An increased atomic number means increased IDID forces, which in turn means increased MP and BP ( the energy required to break them).

79
Q
A

Prelim (4 decks)

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