Module 1: Properties and structure of Matter Flashcards
1
Q
Define Physical properties
A
relate to intrinsic qualities of matter, for example their density.
2
Q
what are physical changes
A
relate to changes in the state of a material but not its chemical structure.
3
Q
what are chemical changes
A
involve a change in chemical structure and composition.
4
Q
what does homogenous mean?
A
Homogenous: Have a uniform composition throughout
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.
6
Q
examples of homogeneous:
A
water, petrol, sugar, aluminium foil
7
Q
examples of heterogeneous:
A
fruit cake, concrete, wood, orange juice
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.
9
Q
Filtration:
Properties + example
A
Physical properties: difference in particle sizes
E.g. sand and sea water
10
Q
Evaporation:
Properties + example
A
Physical properties: Differences in boiling points
E.g. Salt from sea water
11
Q
Distillation:
Properties + example
A
Physical properties: Differences in boiling points of miscible substances
E.g. Alcohol (ethanol) and water in solution
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
13
Q
Decantation:
Properties + example
A
Physical properties: Difference in density
E.g. pouring tea off of tea leaves
14
Q
Separating funnel:
properties + example
A
Physical properties: Difference in immiscible liquids
E.g. Mixture of petrol and water
15
Q
Sedimentation:
properties + example
A
Physical properties: Difference in density
E.g. River sand and gold
16
Q
define miscible
A
Miscible: (of liquids) forming a homogeneous mixture when added together.
17
Q
define Immiscible
A
Immiscible: (of liquids) not forming a homogeneous mixture when mixed.
18
Q
How do you calculate composition?
A
Calculation composition = mass of substance/ total mass x 100
19
Q
Define inorganic substance:
A
Inorganic substances are a group of chemicals that contain no carbon
20
Q
Sulfite ion
A
SO3 ( -2)
21
Q
Sulfate ion
A
SO4 ( -2)
22
Q
Nitrate ion
A
NO3 (-1)
23
Q
Carbonate ion
A
CO3 (-2)
24
Q
Phosphate ion
A
PO4 (-3)
25
Ammonium ion
NH4 (+1
26
Hydroxide ion
OH (-1)
27
what is a covalent bond?
a chemical bond that involves the sharing of electron pairs between atoms
28
What is a polyatomic ion?
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.
29
how do compounds form?
Compounds formed from polyatomic ion combinations and are called polyatomic ionic compounds. But the polyatomic ion behaves as a single unit.
30
Steps in naming (IUPAC)
- 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
31
Physical properties definition + examples
can be observed or measured without changing the composition of matter/ melting point, color, hardness, state of matter, odor, and boiling point
32
Chemical properties definition + examples
a characteristic of a substance that may be observed when it participates in a chemical reaction/ flammability, toxicity, chemical stability, and heat of combustion
33
Pure substance definiton:
that is made up of just one chemical element or compound
34
Impure substance definiton
made of two or more elements or compounds that are not bonded together chemically
35
Elements
only one type of atom
36
Compounds
different atoms in a fixed ratio
37
non-metals
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.
38
Noble gases
Group 18/ full valence shells, and tend to neither gain nor lose electrons
39
transition metals
Group 3-12/ solid at room temperature, except mercury, and have the metallic color and malleability expected of metals
40
alkali metals
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
41
Alkaline earth metals
Group 2/ two valence electrons/slightly harder and less reactive/ rarely found in their elemental form.
42
State at room temp definition
metals
non metals
The form they take at room temp
Solid except mercury
Gases and solids
43
State at room temp defintion
The form they take at room temp
44
Lustre definition
metals
non metals
A gentle sheen or soft glow
Reflect light from its surface and can be polished
No metallic lustre and don’t reflect light
45
conductivity definition
metals
non metals
The degree to which a material conducts electricity
Good conductors of heat and electricity
Poor conductors of heat and electricity
46
ductility definition
metals
non metals
Ability of a material to stretch things
Can be drawn into wires
non-ductile
47
ductility definition
metals
non metals
Ability of a material to stretch things
Can be drawn into wires
non-ductile
48
What are subatomic particles ?
Protons ( + charged)
Neutrons (no charge)
Electrons ( - charged)
49
Where are the subatomic particles found ?
Protons and neutrons found in nucleus. Electrons found orbiting the nucleus
50
what is an atomic number ?
Atomic number= number of protons in an element= number of electrons in a neutral atom
51
what is atomic mass ?
Atomic mass number= number of protons + number of neutrons
52
What are isotopes ?
atoms with the same number of protons, but different numbers of neutrons. They have the same atomic number (same element)
53
When is an isotope stable ?
when it has a stable nucleus and does not emit radiation
54
What is a radioisotope ?
An isotope that has an unstable nucleus and emits radiation
55
When is a nucleus unstable?
- 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
How does a nucleus become stable?
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
What are the types of particles (isotopes)?
- 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
What is the difference between the different types of decay?
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
What is the driving force behind chemical activity?
the atoms want to become ions in order to become stable
60
Spdf numbers e.g. s=? p=?
s=2 p=6 d=10 f=14
61
What could Bohr's model not do?
- 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
What are the amount of electrons in the first four energy levels?
(2,8,18,32)
63
What happens when the third energy level tries to become full?
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
Explain how a flame test works:
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
State of matter at room temperature trends + definition
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
Electronic configurations and atomic radii trends + definition
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
Electronic configurations and atomic radii:
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
Ionisation energy trends + definition
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
Electronegativity trends + definition
measure of tendency of an atom of an element to attract electrons. increases across and then decreases going down
70
Reactivity with water trends + definition
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
Why do atoms bond with other atoms?
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
what can electronegativity help determine
Electronegativity is a key determinant in helping us identify the nature of the bond which may form between two atoms.
73
what is the difference between ionic and covalent bonds
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
ionic bonds definition
Ionic bonds= they take/ give away an electron/as the difference in electronegativity increases an ionic bond is more likely
75
covalent bonds definition
Covalent Bonds= Sharing electrons/ very strong/ when they have similar electronegativities/ two metals
76
how do positive ions form
A positive ion= forms when an atom loses electrons
77
how do negative ions form
A negative ion= forms when an atom gains electrons
78
ionic compound
- form when
- what do they make
- b.p. etc.
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
polyatomic atom
more than one type of atom
80
how are covalent compounds formed
formed by atoms sharing their outer electron.
81
how do you do lewis dot diagrams
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
what is a non polar covalent bond
bonding electrons shared equally between two atoms. No charges on atoms.
83
what is a polar covalent bond
bonding electrons shared unequally between two atoms. Partial charges on atoms.
84
what do electron pairs do in the valance shell
repel each other and are arranged as far away from each other as possible.
85
what are lone pairs
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
how do we use electronegativity to work out bonds
If the electronegativity difference is >1.5, this generally gives ionic bonds, while if it is <1.5, it gives a covalent bond.
87
what happens when the atoms are not the same?
| -in regards to electronegativity
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
what is a intramolecular force
| + examples
Intramolecular forces are the forces involved in chemical bonds. They are within molecules and are relatively strong bonds. E.g. ionic and covalent
89
what is a intermolecular force
| + examples
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
what is dipole dipole
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
what is disperssion
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
what causes temporary/induced polarity
(result of uneven electron distribution within atoms and between neighbouring atoms and molecules).
it is a temporary dipole or charge separation
93
how does molecular weight affect m.p. and b.p.
M.p and b.p increases with molecular weight. This is due to an increase in the strength of the force.
94
what are hydrogen bonds
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
what is an ionic network
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
what is a covalent network
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
what is a covalent molecular
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
what is metallic (chemical structure)
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
define allotrope with an example
An allotrope is when a chemical element comes in many different forms. A common example is Carbon.
100
Graphite
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
Diamond
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
graphene
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
fullerenes
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
carbon nanotubes
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
