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Chemistry > Properties & Structures of Materials > Flashcards

Properties & Structures of Materials Flashcards

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1
Q

Pure substances

A

have a definite and fixed position
not able to be seperated by physical means

elements: one type of atom, no chemical seperation
compounds: two or more elements chemically combined in a definite whole number ratio, chemical seperation possible

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2
Q

Mixture

A

composed of 2 or more different kinds of substances that are physically combined (not chemically combined)

seperated by physical means

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3
Q

Homogenous

A

evenly mixed; uniform composition
called solutions
- air
- seawater
- soft drinks
- brass

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4
Q

Heterogenous

A

non-uniform composition
- concrete
- orange juice
- clay
- wood

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5
Q

Properties of pure substances

A

well-defined and constant

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6
Q

Properties of mixtures

A

vary depending on relative amount of substances that make up the mixture

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7
Q

Physical properties

A
  • mp & bp
  • density
  • solubility
  • electrical and thermal conductivity
  • malleability and ductilability
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8
Q

Chemical properties

A
  • decomposition of heat
  • effect of light
  • reactions with water, acids, bases and oxygen

relates to the ability of a substance to react to form a new substance

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9
Q

Indicators of chemical reactions

A
  • solid precipitate forms
  • gas produced (effervescence)
  • insoluble solid disappears
  • colour change
  • significant change in temperature
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10
Q

Sieving

A

procedure:
2 solids
particle is placed through sieve, larger particles are trapped
property:
particle size

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11
Q

Filtration

A

procedure:
insoluble solids from a liquid
liquid passes through filtrate and solid residue is left behind
property:
particle size & solubility

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12
Q

Evaporation

A

procedure:
retrieve a solid that was dissolved in liquid
heat causes liquid to turn into gas leaving behind solid that was previously dissolved
property:
boiling point

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13
Q

Gravity separation

A

procedure:
2 solids
mixture is typically grounded to fine particle size
agitating (shaking) mixture causes denser particles to settle at bottom and less dense at top
property:
density

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14
Q

Distillation

A

procedure:
solutions with significantly diff bp or a solid from liquid
sample is boiled, solvent changes to vapour, vapour passes down condenser where it is cooled back to a liquid (distillate)
property:
boiling point

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15
Q

Fractional distillation

A

procedure:
liquids with similar bp
components (fractions) with diff bp rise up fractional column to different heights
lowest bp rises highest
highest bp collected at bottom
property:
boiling points

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16
Q

Decantation

A

procedure:
seperate solids from a liquid by pouring off liquid to leave solid behind
property:
density

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17
Q

Magnetic seperation

A

procedure:
mixture is bought into an electric field, particles will either be attracted or repelled
property:
magnetism

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18
Q

Metallic bonding structure

A
  • a lattice of positive ions surrounded by a sea of electrons held by electrostatic attraction
  • valence electrons are held very weakly and are delocalised
  • non-directional bonding
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19
Q

Delocalised

A

Not associated with one particular cation but move around lattice

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20
Q

Metallic substances properties

electrical & thermal conductivity, malleable/ductile, mp/bp, colour

A
  • Good electrical conductor: delocalised electrons can move and conduct charge
  • Good conductor of heat: electrons are mobile so can carry heat (kinetic) energy; vibration of metallic ions contribute to flow of heat through lattice
  • Malleable & ductile: non-directional bonding and delocalised electrons mean metal atoms can move by force without breaking bonds or fracturing
  • High mp & bp: more electrons = higher electrostatic attraction
  • silvery in colour
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21
Q

Ionic Bonding Structure

A
  • strong electrostatic attraction between cation and anion
  • 3D lattice structure
  • every cation is surrounded by 6 anions and vice versa
  • non-directional bonding
  • only movement is vibration as ions are in fixed position
    - unlike charges are adjacent to overpower repulsion from like particles
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22
Q

Coordination number

A

number of ions with opposite charge surrounding the ion

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23
Q

Properties of Ionic Substances

conduction, hardness, brittle, mp/bp, solubility

A
  • doesn’t conduct electricity at solid state: fixed position unable to move and carry charge
  • good conductors of electricity at molten aqueous state: ions are mobile and can carry charge
  • hardness: strong ionic bonds in fixed position
  • brittle: passing over a later of atoms would cause like charge to be next to each other -> repulsion -> fracturing
  • high mp & bp: strong ionic bonds and forces in lattice
  • solubility: varying solubilities; insoluble in non-polar solutes (oil)
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24
Q

Covalent bonding

A

2 non-metals bonded by sharing electrons
- restricted to atoms having similar or high electronegativity as it wants to gain electrons to achieve noble gas configuration
- directional bonds
- held together by strong electrostatic forces between shared electrons and nucleus

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25
Covalent molecular structure and forces
- simple molecules - **intermolecular forces** between molecules are **weak** - **intramolecular forces** between atoms in molecules are **strong** -> electrostatic forces
26
Properties of covalent molecular substances | conduction, soft/weak, mp/bp
- **non-conductors of electricity:** no mobile particles; some are good conductors in aqueous state (acids & bases) - **soft and weak:** weak ***intermolecular*** forces - **low to moderate bp & mp:** only weak ***intermolecular*** forces need to be overcome to melt or boil
27
Covalent network structure
- strong covalent bonds *(bc of strong electronegativity)* - localised electrons - structure is different depending on type
28
Properties of covalent network substances | mp/bp, hardness, brittle, chemical reactivity, solubility
- **very high mp & bp:** strong ***covalent bonds*** - **non-conductors of electricity in solid & liquid state:** no mobile electrons and no ions (exception: graphite) - **extremely hard:** strong ***covalent bonds*** - **brittle:** if some bonds are broken, lattice will be under stress then shatter - **reasonably chemically inert:** all bonding requirements are fulfilled, large amounts of energy needed to rearrange existing bonds - **insoluble in water & most other solvents:** large amounts of energy needed to break bonds
29
Allotropes of Carbon
- diamond - graphite - fullernes - buckyballs - nanotubes
30
Examples of covalent network substances
- Boron - **Carbon** -> graphite & diamond - Silicon - Silicon Carbide - **Silicon dioxide**
31
Diamond
- **Each Carbon has 4C bonded** - tightly bonded electrons, no delocalised -> non-conductor - hard bc of **3D structure** ## Footnote only structure is specific, properties are the same but needed to be compared to graphite
32
Graphite
- **slippery:** 2D structure of flat hexagonal layers are held together by ***weak intermolecular forces*** - **conduct electricity:** each C has 3C bonded but the 4th is delocalised -> ***mobile particles***
33
Law of Conservation of Mass
- atoms are neither created nor destroyed - **mass reactants = mass products** - products only change **chemically**
34
Acid-base reactions
Acid + base -> salt + water
35
Acid-metal reactions
Acid + metal -> salt + hydrogen gas
36
Acid-carbonate reactions
Acid + carbonate -> salt + water + carbon dioxide
37
What are chemical reactions?
Processes in which substances (reactants) are chemically changed into products
38
Mole
a measure of number of particles in a substance
39
molar mass
1 g/mol equal to mass of one mole of a substance
40
the substance does not conduct electricity as a solid, but does as a liquid
ionic
41
when talking about why it doesnt conduct electricity, we have to mention
- where all the electrons are bonded *(e.g. shared in diatomic molecule, bonded to carbon)* - so it's delocalised to carry charge
42
to describe explanation for mp/bp
- forces, bonds and **structure** - how much energy needed to break bonds
43
Octet rule
atoms tend to lose, gain or share electrons until they are surrounded by 8 valence electrons
44
VSPER
- due to electron-electron repulsion, bonds and lone pairs are arranged as far apart as possible - accounts for molecular shape *- electrons in a bond don't repel as much as lone pairs* | *valence shell electron pair repulsion*
45
Linear
- 2 domains - 0 lone pairs - bond angle = 180
46
Triangular planar
- 3 domains - 0 lone pairs - bond angle = 120
47
Tetrahedral
- 4 domains - 0 lone pairs - bond angle = 109.5
48
Pyramidal
- 3 domains - 1 lone pair - bond angle = <109.5
49
V-shaped (bent)
- 2 domains - 2 lone pairs - bond angle = <109.5
50
What are exceptions to molecular covalents?
- BeCl2 - BCl3
51
Electronegativity
electron attracting power of an atom involved in a bond measured 0-4
52
Non-polar covalent bond
- 2 identical atoms covalently bonded
53
Polar covalent bond
- 2 atoms with different electronegativity - electron pair is **not shared equally** - more with the higher electronegativity - they have a bond dipole - **uneven charge distribution**
54
Delta
- **delta positive** -> partial positive charge -> less electronegativity - **delta negative** -> partial negative charge -> more electronegativity
55
Extreme polar bonds
ionic bonds
56
Strong bonds
- covalent - ionic - metallic
57
Dispersion forces
occur in all particles
58
Weak intermolecular forces
- Hydrogen bonds - Dipole-dipole bonds - Dispersion forces
59
How are dispersion forces formed
- electron random movement - cause **temporary dipoles** - this induces the opposite dipole of the nearby molecule (**induced dipole**) - attractive force that exists between two temporary dipoles is **dispersion force**
60
What are Dispersion forces
attractive force that exists between 2 temporary dipoles | in all molecules
61
How do dispersion forces form?
- random motion of electrons results in temporary dipoles - causing surrounding molecules to form induced dipoles
62
Factors affecting dispersion forces
- **more electrons** - (**larger atomic radius**) - **surface area**
63
How does atomic radius affect dispersion forces
extra electrons has **greater volume** of space to move around greater chance of being asymmetrically arranged
64
How does surface area affect dispersion forces
larger surface area = stronger attractions e.g. linear
65
What are dipole-dipole forces
attractive forces between 2 permanent dipoles of polar molecules
66
What effect do dipole-dipole forces have on mp and bp and solubility
higher mp and bp will be soluble in polar solvents
67
Strongest dipole-dipole forces include
N,O,F - **strong electronegativity** - **small radius** -> partial charge is in small area of space -> high charge density at end of each molecule -> more polar
68
Hydrogen Bonds
attraction between polar molecules containing OH, NH or FH groups *hydrogen is bonded to lone pairs*
69
Hydrogen bonds solubility and mp/bp
Much higher mp and bp extremely soluble
70
Ion-dipole forces
attraction between polar solvent water molecules and dissolved ion *higher charge of ion means more attraction*
71
Explain solubility in terms of ion-dipole forces
- ion-dipole forces are much stronger than ion bonds - ion dipole interactions **replace ionic bonds** - more strong an ionic bond is, less solubility
72
Changing phase | and factors
between solid to liquid or liquid to gar - more **intermolecular force** = higher mp and bp - increase in **size** = increase in mp/bp
73
Solubility
solute and solvent can only form a solution if both molecules are **comparable in strength** - *like dissolves like* - non-polar dissolves non-polar, polar dissolves polar
74
Define Equillibrium vapour pressure | and factor (temp)
measure of a tendency of a substance to evaporate (kPa) - higher temp = higher kPa -> *bc more kinetic energy that better enable molecules to escape molecular forces*
75
What substances have higher kPa
the ones with the weakest forces, they evaporate the easiest
76
Unique properties of water
- high bp/mp - denser in liquid phase - high surface tension
77
Why does water have high mp and bp
able to form **extensive hydrogen bonding** with neighbouring molecules
78
Why is water denser in liquid phase
- in ice, the lattice (diamond ish structure) allow water to form 4 hydrogen bonds with 4 neighbouring water molecules - this takes up **more space** (**less dense**) than the ***random arrangments*** of molecules in liquid water
79
Define surface tension
liquids tendency to resist any increase in it's surface area
80
Why does water have a high surface tension
- **strong intermolecular forces** - at the surface there is an **imbalance** of these forces, causing molecules to be pulled inwards into the bulk of the liquid - causes water to form spherical droplets -> that's the shape which allows least surface area to volume ratio

Chemistry (4 decks)

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