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Chemistry preliminary 2020 > Module 1: Properties and structure of Matter > Flashcards

Module 1: Properties and structure of Matter Flashcards

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1
Q

Define Physical properties

A

relate to intrinsic qualities of matter, for example their density.

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2
Q

what are physical changes

A

relate to changes in the state of a material but not its chemical structure.

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3
Q

what are chemical changes

A

involve a change in chemical structure and composition.

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4
Q

what does homogenous mean?

A

Homogenous: Have a uniform composition throughout

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5
Q

what does heterogenous mean?

A

Have non-uniform composition throughout. This means we can recognise different particles in the mixture as they have very different properties.

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6
Q

examples of homogeneous:

A

water, petrol, sugar, aluminium foil

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7
Q

examples of heterogeneous:

A

fruit cake, concrete, wood, orange juice

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8
Q

What is the difference between homogenous and heterogenous mixtures?

A

A Homogenous mixture has a uniform composition throughout such as in petrol or pure water. Whereas a heterogeneous mixture refers ti having a non-uniform composition resulting in one being able to recognise small pieces of the material that are different from other pieces. For example in fruit cake and orange juice.

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9
Q

Filtration:

Properties + example

A

Physical properties: difference in particle sizes

E.g. sand and sea water

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10
Q

Evaporation:

Properties + example

A

Physical properties: Differences in boiling points

E.g. Salt from sea water

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11
Q

Distillation:

Properties + example

A

Physical properties: Differences in boiling points of miscible substances
E.g. Alcohol (ethanol) and water in solution

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12
Q

Fractional Distillation:

Properties + example

A

Physical properties: Small difference in boiling points

E.g. Removing substances such as kerosene, gasoline, diesel from crude oil

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13
Q

Decantation:

Properties + example

A

Physical properties: Difference in density

E.g. pouring tea off of tea leaves

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14
Q

Separating funnel:

properties + example

A

Physical properties: Difference in immiscible liquids

E.g. Mixture of petrol and water

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15
Q

Sedimentation:

properties + example

A

Physical properties: Difference in density

E.g. River sand and gold

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16
Q

define miscible

A

Miscible: (of liquids) forming a homogeneous mixture when added together.

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17
Q

define Immiscible

A

Immiscible: (of liquids) not forming a homogeneous mixture when mixed.

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18
Q

How do you calculate composition?

A

Calculation composition = mass of substance/ total mass x 100

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19
Q

Define inorganic substance:

A

Inorganic substances are a group of chemicals that contain no carbon

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20
Q

Sulfite ion

A

SO3 ( -2)

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21
Q

Sulfate ion

A

SO4 ( -2)

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22
Q

Nitrate ion

A

NO3 (-1)

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23
Q

Carbonate ion

A

CO3 (-2)

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24
Q

Phosphate ion

A

PO4 (-3)

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25
Q

Ammonium ion

A

NH4 (+1

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26
Q

Hydroxide ion

A

OH (-1)

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27
Q

what is a covalent bond?

A

a chemical bond that involves the sharing of electron pairs between atoms

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28
Q

What is a polyatomic ion?

A

Polyatomic ions are covalently bonded groups of atoms with a positive or negative charge caused by the formation of an ionic bond with another ion.

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29
Q

how do compounds form?

A

Compounds formed from polyatomic ion combinations and are called polyatomic ionic compounds. But the polyatomic ion behaves as a single unit.

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30
Q

Steps in naming (IUPAC)

A
  • Always name the metal first. It keeps it whole name, eg magnesium, sodium, copper, etc.
  • If there is 1 non-metal, its ending is changed to ‘ide’. -Chlorine/chloride, oxygen/oxide, bromine/bromide.
  • If there are 2 non-metals with a metal, one of which is oxygen, the other non-metal gets an ‘ate’ ending, e.g. nitrogen + oxygen= nitrate, carbon + oxygen = carbonate and so on
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31
Q

Physical properties definition + examples

A

can be observed or measured without changing the composition of matter/ melting point, color, hardness, state of matter, odor, and boiling point

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32
Q

Chemical properties definition + examples

A

a characteristic of a substance that may be observed when it participates in a chemical reaction/ flammability, toxicity, chemical stability, and heat of combustion

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33
Q

Pure substance definiton:

A

that is made up of just one chemical element or compound

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34
Q

Impure substance definiton

A

made of two or more elements or compounds that are not bonded together chemically

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35
Q

Elements

A

only one type of atom

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36
Q

Compounds

A

different atoms in a fixed ratio

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37
Q

non-metals

A

All elements on the right besides group 18/ Many of the nonmetals are gaseous, and all are notable for their tendency to gain electrons and fill their valence shells.

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38
Q

Noble gases

A

Group 18/ full valence shells, and tend to neither gain nor lose electrons

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39
Q

transition metals

A

Group 3-12/ solid at room temperature, except mercury, and have the metallic color and malleability expected of metals

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40
Q

alkali metals

A

Group 1/ highly reactive metals/ only one electron in their valence shell, which is easily donated to another atom in chemical reactions/ explosive reactivity in both air and water/ rarely found in their elemental form in nature

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41
Q

Alkaline earth metals

A

Group 2/ two valence electrons/slightly harder and less reactive/ rarely found in their elemental form.

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42
Q

State at room temp definition
metals
non metals

A

The form they take at room temp
Solid except mercury
Gases and solids

43
Q

State at room temp defintion

A

The form they take at room temp

44
Q

Lustre definition
metals
non metals

A

A gentle sheen or soft glow
Reflect light from its surface and can be polished
No metallic lustre and don’t reflect light

45
Q

conductivity definition
metals
non metals

A

The degree to which a material conducts electricity
Good conductors of heat and electricity
Poor conductors of heat and electricity

46
Q

ductility definition
metals
non metals

A

Ability of a material to stretch things
Can be drawn into wires
non-ductile

47
Q

ductility definition
metals
non metals

A

Ability of a material to stretch things
Can be drawn into wires
non-ductile

48
Q

What are subatomic particles ?

A

Protons ( + charged)
Neutrons (no charge)
Electrons ( - charged)

49
Q

Where are the subatomic particles found ?

A

Protons and neutrons found in nucleus. Electrons found orbiting the nucleus

50
Q

what is an atomic number ?

A

Atomic number= number of protons in an element= number of electrons in a neutral atom

51
Q

what is atomic mass ?

A

Atomic mass number= number of protons + number of neutrons

52
Q

What are isotopes ?

A

atoms with the same number of protons, but different numbers of neutrons. They have the same atomic number (same element)

53
Q

When is an isotope stable ?

A

when it has a stable nucleus and does not emit radiation

54
Q

What is a radioisotope ?

A

An isotope that has an unstable nucleus and emits radiation

55
Q

When is a nucleus unstable?

A
  • too many neutrons or protons
  • more protons increases repulsive forces, causing instability
  • more neutrons increase attractive forces, causing instability
  • All nuclei with 83+ protons are unstable as the repulsive forces are too large
56
Q

How does a nucleus become stable?

A

The nucleus will try to ‘fix’ itself to become stable and in the process will emit different particles depending on how the nucleus becomes stable.

57
Q

What are the types of particles (isotopes)?

A
  • Alpha particle: 2 protons, 2 neutrons (helium nucleus). They are relatively heavy, positively charged and have a low penetrating power. Blocked by paper and skin.
  • Beta particle: electron. Lighter, negatively charged and have a greater penetrating power. Blocked by aluminium.
  • Gamma radiation: energy, no charge, and are extremely penetrating. Blocked by lead and concrete.
58
Q

What is the difference between the different types of decay?

A

Alpha decay: atomic number above 83, quick way to lose mass. Atomic number changes by 2 and atomic mass changes by 4

Beta decay: when a nucleus has too many neutrons. A neutron is converted into a proton and a beta particle. Change in atomic mass by +1.

Gamma decay: occurs with many nuclear decay reactions. When the nuclear particles rearrange themselves and excess energy is emitted as high energy electromagnetic rays.

59
Q

What is the driving force behind chemical activity?

A

the atoms want to become ions in order to become stable

60
Q

Spdf numbers e.g. s=? p=?

A

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

61
Q

What could Bohr’s model not do?

A
  • accurately predict emission spectra of atoms with more than one electron
  • explain why electron shells can only hold 2n(squared) electrons
  • Explain why the fourth shell accepts 2 electrons before the third shell is full
62
Q

What are the amount of electrons in the first four energy levels?

A

(2,8,18,32)

63
Q

What happens when the third energy level tries to become full?

A

he 3rd shell does not immediately fill with 18 electrons. It fills with 8, before filling the 4th shell with 2, then returning to complete filling up the 3rd shell. Therefore, the rules only apply up to element 18.

64
Q

Explain how a flame test works:

A

When we heat atoms it causes electrons to become ‘excited’ and they jump into a higher energy level. When they drop back to their ground state, they release light energy.

The more energy released, the shorter the wavelength of light emitted, therefore changing the colour of light emitted. When the emitted light is split into its components, it is found that it occurs at specific wavelengths.

Each of the energy sublevels have a different energy n atoms of different elements. Hence, why each element has a unique emission spectrum.

65
Q

State of matter at room temperature trends + definition

A

State of matter at room temperature: (Melting and boiling point trends)
The stronger the attracting force between atoms, the higher the atoms boiling point will be.
Solid:
Liquid:
Gas:

66
Q

Electronic configurations and atomic radii trends + definition

A

Electronic configurations and atomic radii: decreases across a period= the increasing number of protons attracts the electrons. Increases as we move down a periodic table= outer electrons are further away from the protons.

67
Q

Electronic configurations and atomic radii:

A

Electronic configurations and atomic radii: decreases across a period= the increasing number of protons attracts the electrons. Increases as we move down a periodic table= outer electrons are further away from the protons.

68
Q

Ionisation energy trends + definition

A

the amount of energy required to move an electron from an atom in its gaseous state. First ionisation energy refers to the energy needed to remove the first electron. Increases along periods= increased core charge. Decreases down a group= easier to remove an electron

69
Q

Electronegativity trends + definition

A

measure of tendency of an atom of an element to attract electrons. increases across and then decreases going down

70
Q

Reactivity with water trends + definition

A

Group 1 and 2, increases from top to bottom. The larger the atomic radius the easier it is to lose an electron. Reactivity decreases left to right= more electrons in the outer shell.

71
Q

Why do atoms bond with other atoms?

A

Atoms bond with other atoms to become stable. In order to become stable they try and get a full outer shell. They do this by either forming ionic or covalent bonds in which they either take, give or share electrons.

72
Q

what can electronegativity help determine

A

Electronegativity is a key determinant in helping us identify the nature of the bond which may form between two atoms.

73
Q

what is the difference between ionic and covalent bonds

A

Ionic bonds= they take/ give away an electron/as the difference in electronegativity increases an ionic bond is more likely

Covalent Bonds= Sharing electrons/ very strong/ when they have similar electronegativities/ two metals

74
Q

ionic bonds definition

A

Ionic bonds= they take/ give away an electron/as the difference in electronegativity increases an ionic bond is more likely

75
Q

covalent bonds definition

A

Covalent Bonds= Sharing electrons/ very strong/ when they have similar electronegativities/ two metals

76
Q

how do positive ions form

A

A positive ion= forms when an atom loses electrons

77
Q

how do negative ions form

A

A negative ion= forms when an atom gains electrons

78
Q

ionic compound

  • form when
  • what do they make
  • b.p. etc.
A

form when a + ion is attracted to a - ion. This is known as electrostatic attraction. These are ionic bonds.Must be a metal and non-metal. Have a high boiling point and are soluble.

79
Q

polyatomic atom

A

more than one type of atom

80
Q

how are covalent compounds formed

A

formed by atoms sharing their outer electron.

81
Q

how do you do lewis dot diagrams

A

Lewis Dot diagrams:
Calculate valence electrons
Place them in appropriate positions
Place a dash to make a single bond
Arrange so the outer electrons have a full outer shell
If the central atom does not have 4 electron pairs, make a double or triple bond

82
Q

what is a non polar covalent bond

A

bonding electrons shared equally between two atoms. No charges on atoms.

83
Q

what is a polar covalent bond

A

bonding electrons shared unequally between two atoms. Partial charges on atoms.

84
Q

what do electron pairs do in the valance shell

A

repel each other and are arranged as far away from each other as possible.

85
Q

what are lone pairs

A

are electrons that are not involved in the bonding of atoms. They are treated in the same way as the bonding pairs when determining the shape of a molecule.

86
Q

how do we use electronegativity to work out bonds

A

If the electronegativity difference is >1.5, this generally gives ionic bonds, while if it is <1.5, it gives a covalent bond.

87
Q

what happens when the atoms are not the same?

-in regards to electronegativity

A

When the atoms are not the same, the more electronegative element will hold the electrons closer. These bonds are said to have a dipole (positive end and negative end). The more electronegative element becomes the negative end.

88
Q

what is a intramolecular force

+ examples

A

Intramolecular forces are the forces involved in chemical bonds. They are within molecules and are relatively strong bonds. E.g. ionic and covalent

89
Q

what is a intermolecular force

+ examples

A

Intermolecular forces are the forces of attraction between molecules, they are weaker. They are physical bonds and are the ones affected by physical changes of state.
-dipole dipole
disperssion
hydrogen bonds

90
Q

what is dipole dipole

A

occurs between molecules which are polar. Polar bonds are the result of permanent dipole-dipole interactions. They have the effect of raising m.p. And b.p. As they are quite strong forces.

91
Q

what is disperssion

A

Molecules that are neither ionic, nor polar covalent still need something holding them together otherwise they would be gases.
They result from temporary induces polarity within the molecules The shared electrons may be closer to one atom than the other giving them an induced polarity.
They are the result of temporary dipole dipole interactions. They are weaker.
Larger atoms display stronger dispersion forces than smaller atoms.
M.p and b.p increases with molecular weight. This is due to an increase in the strength of the force.

92
Q

what causes temporary/induced polarity

A

(result of uneven electron distribution within atoms and between neighbouring atoms and molecules).
it is a temporary dipole or charge separation

93
Q

how does molecular weight affect m.p. and b.p.

A

M.p and b.p increases with molecular weight. This is due to an increase in the strength of the force.

94
Q

what are hydrogen bonds

A

Stronger than usual.
It occurs in oxygen, nitrogen or Fluorine.
Hydrogen becomes slightly positive, attracting the non-bonding electron pair of other atoms through electrostatic forces.
Seen in DNA and Protein.
M.p. and B.p. is high.

95
Q

what is an ionic network

A

positive and minus charges attracted to one another, hold in lattice shape. E.g. NaCl
They are held together by electrostatic attraction. Very high temp needed to break strong bonds. Low conductivity as ions are tightly bound and unable to move towards change. They are hard and brittle.

96
Q

what is a covalent network

A

Covalent bonding extends through the crystal in diamond. E.g. SiO2
Held together by covalent bonds. Very high temp needed to break strong covalent bonds. Low conductivity as no ions and no electrons are free to move. They are hard but no malleability.

97
Q

what is a covalent molecular

A

Molecules on their own with low intermolecular forces. E.g. Br2O2
Held together by covalent (intra) and dispersion/dipole dipole (inter). Low mp and bp as not much energy needed to break intermolecular bonds. Low conductivity. Soft/ low hardness/ low malleability

98
Q

what is metallic (chemical structure)

A

Orderly 30 array of positive ions held together by mobile ‘sea’ of delocated electrons. Valence electrons break away leaving positive ions and more randomly. E.g. Copper
Held together by electrostatic. M.p. and B.p. depend on metal. High conductivity as electrons move freely through the lattice. Malleable.

99
Q

define allotrope with an example

A

An allotrope is when a chemical element comes in many different forms. A common example is Carbon.

100
Q

Graphite

A

2D covalent network substance
Each atom is covalently bonded to three others- forms flat hexagonal sheet
Spare electron becomes delocalised and holds the sheets together through weak intermolecular forces – can slide over one another without disrupting bonds
High electrical conductivity due to delocalised electrons § High thermal conductivity- thermal energy can pass through covalent bonds
Soft /Slippery – sheets can slide over one another – dry lubricant/ pencils
Brittle – covalent bonds will shatter (not bend) under excessive pressure
High melting point- high amount of energy needed to break covalent bonds

101
Q

Diamond

A

3D covalent network substance
Each atom is covalently bonded to four other atoms – tetrahedral lattice
Very strong due to stable and rigid bonds, able to resist large amounts of force from many directions – used as a cutting tool
Atoms within the structure cannot move out of their places, without breaking the bonds, cannot bend will only shatter – make it brittle, will shatter under excessive force
High melting point- due to many strong bonds needing high energy to break
High thermal conductivity – atoms are close to one another
Low electrical conductivity – no free charged particles can move/ carry current

102
Q

graphene

A

Single sheet of graphite
High melting point, high electrical conductivity, high thermal conductivity – like graphite
Easily moulded and very flexible – because it is only one atom thick
High tensile strength- can be spun into wires (unlike graphite where the sheets will just slide over one another)

103
Q

fullerenes

A

First discovered when man made in 1985, but also occur naturally
C60 is a hexagonal based sphere arrangement – exists as discrete molecules
High melting point- high amount of energy needed to break bonds
High tensile strength- many covalent bonds can resist force from many directions
Electrical insulator – electrons cannot pass current because they cannot move through entire substance, since each fullerene is a separate molecule
Can act as superconductors is metal is added to them

104
Q

carbon nanotubes

A

Synthetic, long hollow tube
High tensile strength due to tube-like structure – higher than graphene because they can resist and disperse pressure from more directions
Conducts electricity – delocalised electrons can travel the length of the tube
Ductile- can be spun into wires

Chemistry preliminary 2020 (5 decks)

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