Elements of Life Flashcards
1
Q
What is the atomic number?
A
The number of protons in the nucleus of an atom
2
Q
What is the mass number?
A
The total number of protons and nuetrons in the nucleus of an atom
3
Q
What is relative atomic mass (Ar)?
A
The mass of one atom of an element relative to 1/12 the mass of carbon-12
Is an average of relative isotopic masses, taking into account abundance
4
Q
What do the different numbers on the nuclear symbol (of an element etc.) tell you?
A
No. protons = Atomic no. (Bottom no.)
No. electrons = Atomic no. (Bottom no.)
No. neutrons = Mass no. - atomic no.
(Top no. - bottom no.)
Mass no. (Ar) = Top no.
5
Q
How are models of the atom made + updated?
A
Tested using experimental investigations
Are revised when observations are made that aren’t predicted by model
6
Q
What were the different steps/models in the development of the current atomic model?
A
- Dalton model - simple ‘billiard ball’. Particles cannot be divided, created, or destroyed. Are unique.
- ‘Plum pudding’ model - electrons embedded in sea of positive charge. Discovered by firing cathode rays (electrons) in air - discovered e-. Introduced idea that atoms made of smaller particles
- Nuclear model - Geiger-Marsden experiment showed some alpha particles deflected at large angles by small, dense area of positive charge surrounded by e-
- Bohr model - Evidence from atomic spectra + patterns of ionisation enthaply. e- arranged in shells - ‘planetry model’
7
Q
What is nuclear fission?
A
The splitting of a large, unstable isotope triggered by bombarding it with smaller, high-speed particles (usually neutrons)
8
Q
What conditions are needed for nuclear fission?
Why?
A
High temps and/or pressure to provide the energy needed to overcome the repulsion between the 2 positive nuclei
9
Q
What is the nuclear symbol for a neutron?
A
10n
Except with the 1 and 0 in line above each other…
10
Q
Write a nuclear equation for the fusion of a 11H nucleus with a 21H nucleus
A
21H + 31H → 42He + 10n
- Ignore the bit in the picture about what reaction it is…*
- Also neutrons only need to be included if numbers don’t add up*
11
Q
What is the general formula for calculating Ar?
A
(% abundance of x X isotopic mass of x) + (% abundance of y X isotopic mass of y)
/ 100%
12
Q
The Ar of potassium is 39.1. Calculate the relative abundance of 39K and 41K
A
Make 1 isotope abundance x so the other = 100-x
(39x + 41(100-x))/100% = 39.1
Multiply both sides by 100 + multiply out brackets
39x + 4100 - 41x = 3910
-2x = -190, x =95
39K = 95%, 40K = 5%
13
Q
What are isotopes?
A
Atoms of the same element with a different number of neutrons
This causes mass number to be different
Their relative abundances are used to calculate Ar
14
Q
What is Mr?
A
Relative molecular mass
The ratio of the average mass of one molecule of an element or compound to one twelfth of the mass of an atom of carbon-12.
(Ar but for molecules… (not elements))
15
Q
What is the Avogadro constant (NA)?
A
The number of atoms/molecules in 1 mole of a substance
16
Q
What does quantised mean?
A
Energy that can only take particular values (known as quanta)
17
Q
What is the ground state?
A
The lowest energy level that an electron can occupy
18
Q
What is a photon?
A
Quanta of energy in the form of electromagnetic radiation
19
Q
Breifly describe Bohr’s model of the atom
A
Electrons in an atom occupy discrete, quantised energy levels/shells
Electrons in an energy level have a specific amount of energy
Hence the energy of the electron is said to quantised
20
Q
What property does light have?
What does this mean?
A
Wave-particle duality
Means it can behave like both a wave and a particle…
21
Q
What properties does light have the mean it can be described as a particle?
A
Made up of ‘tiny packets of energy’ called photons
The energy of a photon corresponds to its position in the EM spectrum
Increased freq. = increased energy + decreased wavelength
22
Q
What equation links the wave + particle models of light?
A
ΔE = hv
ΔE = energy of photon (J) h = Planck's Constant v = frequence (Hz/s<sup>-1</sup>)
23
Q
What equation expalins the wave properties of light?
A
c = vλ
c= speed of light (ms<sup>-1</sup>) v = frequency (Hz/s<sup>-1</sup>) λ = wavelength (m)
24
Q
Describe the appearance of an emission spectrum
A
Consists of coloured lines on a black background
The lines become closer at higher frequencies
There are several series of lines (although some may fall outside visible part of spectrum)
25
What is spectroscopy?
The study of how light and matter interact
Uses IR, visible, and UV light
26
Explain the formation of an emission spectrum
* **Electrons** in the **ground state absorb energy**
* This promotes them to a **higher energy level** - **excited state**
* Electrons then **drop back down** to lower energy levels. The **energy lost** (ΔE) us **emitted as a photon** of light
* The **frequency** of the photon is **related** to the **energy** lost by **ΔE = hv**
* **Different energy gaps** produce **photons of different frequencies**
* This produces **different coloured bands** on the emission spectrum
27
Why can emission/absorption spectra be used to identify different atoms from a compound/mixture?
Because each element has a **unique configuration of electrons**, therefore has a unqiue emission/absorption spectrum
The **energy levels of the electrons are discrete + quantised** means **only certain freqs. emitted/absorbed** - it's **not continuous**
28
What are flame tests?
Used to identify the presence of specific metals in a sample
Different metals give different coloured flames depending on their emission spectra
29
What is flame colour?
The light emitted by metal ions when a vaporised metal salt is heated up in a flame
30
What colour flame does Li+ give?
*Bright* red
31
What colour flame does Na+ give?
Yellow
32
What colour flame does K+ give?
Lilac
33
What colour flame does Ca2+ give?
*Brick* red
34
What colour flame does Ba2+ give?
*Apple* green
35
What colour flame does Cu2+ give?
Blue-green
36
Describe the appearance of an absorption spectrum
If white light is passed through a sample of vaporised atoms, an absorption spectrum is seen
Shown by black lines on a rainbow background (showing all colours of visible light)
37
How are atomic absorption spectra formed?
* Electrons in the **ground state absorb photons** of light
* The energy from these photons causes the **electrons to be excited to higher** energy levels
* The electrons **drop back down** to the ground state and a photon/**light is emitted**
* The **energy of this photon** is **related** to the frequency/**energy of light initally** absorbed as **ΔE = hv**
* **Light** of the frequency **doesn't pass through the sample** (as it's absorbed) so a **black line** is seen in the spectrum
38
What are the similarities between emission and absorption spectra?
For a given element, **lines** appear at the **same frequency**
**Lines converge** at a **higher frequency**
**Several series of lines** are seen
39
What are the differences between atomic emission and absorption spectra?
**Emission** spectra show **coloured lines on a black background**
**Absorption** spectra show **black lines on a coloured background**
40
Why do the lines of emission/absorption spectra get closer together at higher frequencies?
**Higher frequency lines** are caused by **translations of electrons with large ΔE values**
These are **produced** from translations from **higher energy levels**
**Higher energy levels are much closer** together than lower energy levels
Translations from **adjacent energy levels** will have **similar ΔE** values and hence **produce light of similar frequencies**
41
Why are several series of lines seen on emission/absorption spectra?
**Lines** are **produced** when **electrons drop to a lower energy level**
Different series of lines are produced by electrons **dropping to different ground states**/electron energy levels
42
What is the principle quantum number?
**Shell**
Given as n (i.e. 1,2,3 etc. the number before the letter...)
The higher the value, the higher the energy
43
What are shell divided in to?
**Sub-shells**
Labelled *s*, *p*, *d*, and *f*
44
What is each sub-shell divided into?
What are its properties?
**Atomic orbitals**
Each can **hold max of 2 electrons**
These electrons must have **opopsite** (or paired) **spinds**
Represented by boxes. Arrows drawn in them represent electrons
45
How many orbitals does the s sub-shell contain?
1 s-orbital
46
Summarise the way in which electrons are organised in atoms, starting with the largest grouping.
Shell/PQN
Sub-shells
Atomic orbitals
47
Draw the different shapes of the p orbital
px-orbital
py-orbital
pz-orbital
48
How many orbitals does the p sub-shell have?
3
49
How many orbitals does the d sub-shell have?
5
50
How many orbitals does the f sub-shell have?
7
51
What are the rules that determine the distribution of electrons in atomic orbitals?
The orbitals are **filled in order of** **increasing energy**
Where there is **more than one orbital at the same energy**, the orbitals are **first occupied by a single electron**. **When each orbital is singly occupied**, the **electrons pair up in the orbitals**
Electrons in **singly occupied** orbitals have **parallel spins**
Electrons in **doubly occupied** orbitals have **opposite** (paired) **spins**
52
What are the 2 ways of representing electron distribution?
By writing out the electronic configuration in full
e.g. 1s22s22p5
By drawing the electronic configuration in boxes...
*(see picture)*
53
Draw a diagram to show the shape of the s-orbital
It's a ball shape
54
On the periodic table what is a period?
A row
55
What is periodicity?
The **occurence of a regular pattern** in a property as you go across a **period**
The regular pattern is also **repeated in other periods**
56
How are elements in the periodic table arranged?
How did they used to be arranged?
Arranged by atomic number (no. protons)
Used to be arranged by Ar
57
What trend do melting/boiling points follow across a period?
(e.g. period 3)
**Melting point increases then decreases across the period**
This is because the **metals** on the left-hand side of a period are **metalically bonded** so have higher melting points due to the **deloclaised electrons** between nuclei. The **further across** the period, the **more electrons** and the **more positive the nucleus becomes**, so the stronger the **bonds**.
**Silicone** has a **high melting point** because it is a **giant covalent structure** which requires a lot of energy to break
The **remaining non-metals are simple molecules**. They are only held together by **weak intermolecular forces** (e.g. id-id). The **melt** these molecule you **don't need to break the strong covalent bonds**, only the weak intermolecular bonds.
58
What is first ionisation enthalpy?
The **energy** needed to **remove one electron from each atom in one mole** of **isolated gaseous atoms** of an element
59
What is the general equation for first ionisation enthalpy?
X(g) → X+(g) + e-
*Remember state symbols!*
60
What is the trend in first ionisation enthalpies as you go across a period?
Why?
**First ionisation enthalpy increases across a period**.
**Group 0** elements have the **highest values** because they have **full outer shells**, making it difficult to remove an electron
First ionisation enthalpy is **lowest for Group 1** elements because they have only **1 outer shell electron** which is relatively easy to remove/ionise
First ionisation enthalpy **increases across a period** because the **number of protons in the nucleus increases**, meaning the **electrons are more strongly attracted** to the nucleus so are harder to remove
61
What is the trend in first ionisation enthalpy down a group?
Why?
First ionisation enthalpy **decreases** down a group
This is because **electrons** are in shells that are further away from the nucleus, thus the **attraction between the two is less** (**electron shielding**)
62
What is the trend for atomic radii across a period?
Decreases due to the increased number of protons
This means there is greater attraction between the outer electrons and the nucleus
63
What is a closed shell?
When all sub-shells are fully occupied by electrons
64
Why are s-block elements more reactive than p-block elements?
Because the formation of M+ or M2+ ions only requires input of energy equivalent to the first/second ionisation enthalpy.
For p-block elements greater input of energy is needed to lose eletrons due to the greater electron affinity as a result of a more positive nucleus
65
What is a dative covalent bond?
A type of covalent bond in which **both electrons come from the same atom**
Show by **arrow** pointing **away** from **donor**
66
What is a lone pair?
A **pair of electrons** in the outer shell of an atom that are **not involved in bonding**
67
What is ionic bonding?
Bond formed between **metal + non-metal atom**
**Metal** transfers/**donates electron(s)** to non-metal atom
This results in **formation of charged ions**, often with **full outer shells**. This makes them particularly **stable**
68
Draw a dot-and-cross diagram to show the formation of NaCl from Na and Cl
Note: You only need to show outer shells + don't have to draw circles
(Also don't need the arrow, should be written like an equation)
69
How are individual ions held together to form ionic compounds?
**Cations + anions** produced by ionic bonding held together by **electrostatic attraction** between each other
Results in the formation of a **giant ionic lattice**
70
What is covalent bonding?
Bonding that occurs between 2 **non-metal atoms**
Formed by the atoms **sharing one or more pairs of electrons**
(If 2 pairs shared, double bond formed, etc.)
71
Draw diagrams to show the different ways in which dative covalent bonds can be represented
Can be represented by:
Dot-and-cross diagrams
Arrow (in structural formulae)
72
How are the atoms in a simle covalent bonds?
There is an **electrostatic attraction** between the **positive nuclei** of the 2 atoms and the **shared pair of negative electrons** in the bond
This is **greater than the repulsion between the 2 nuclei**
73
What is electron pair repulsion theory?
(AKA VSEPR - Valance Shell Electron Pair Repulsion)
States that the shape adopted by a simple molecule is that which keeps repulsive forces to a minimum.
All bond angles must add up to = 360º
74
Describe/explain how electron pair repulsion determines the shape of molecules
**Electron pairs/groups repel each other**
They will **arrange** themselves to get **as far apart as possible**
State the **no. total pairs** **of electrons**
State the **no. bondind pairs/groups** of electrons
(if applicable) state the **no. lone pairs**
(if applicable) **lone pairs repel more** than bonding pairs (**decrease** bond angle by **2.5º** each)
This creates the **shape** [...] with the **angle**(s) [...]
75
When describing/explaining electron pair repulsion in double/triple bonds, how should you describe electrons?
As **groups** not pairs
76
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Linear
180º
2 pairs of electrons (all bonding)
77
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Triangular planar
120º
3 pairs of electrons (all bonding pairs)
78
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Tetrahedral
109.5º
4 pairs of electrons (all bonding pairs)
79
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Trigonal Bipyrimidal
90º and 120º
5 pairs of electrons (all bonding pairs)
80
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Octahedral
90º
6 electron pairs (all bonding)
81
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Pyrimidal
107º
4 pairs of electrons
3 bonding pairs
1 lone pair
82
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Bent (or V-shaped)
104.5º
4 electron pairs
2 bonding pairs
2 lone pairs
83
Give the shape and bond angle of the molecule shown in the picture
State how many bonding pairs/pairs of electrons there are in the molecule
Bent (or V-shaped)
118º
3 electron groups
2 bonding groups
1 lone pair
84
How do double/triple bonds affect the number of bonding electrons?
Can be thought of as a **single group of electrons**
e.g. CO2 has double 2 double bonds (4 electron pairs in total) which can be thought of as 2 electron groups
85
What type of structure do ionic bonds have?
Always giant ionic (lattice)
86
What type of structure do covalent bonds have?
Either **simple molecular** or **giant covalent network**
87
What type of structure do metallic bonds have?
Always giant metallic lattice
88
Describe the structure of a giant ionic lattice
Has a **regular repeating pattern** of **postivitely and negatively charged ions** in all **3 dimensions**
The **attraction** between these **oppositely charged ions outweighs the repulsion** between ions with the same charge becayse the **oppositely charged ions are closer**
89
What are the characteristic properties of giant ionic lattices?
**High melting point** because of **strong electrostatic attractions** between ions
**Often soluble in water** (due to **charges** of ions)
**Conduct electricity when molten/in solution** as **charged ions able to move** in response to voltage
90
Describe the structure/bonding in simple molecular covalent bonding
**Strong covalent bonds within molecules** (between atoms) (**strong intramolecular** bonds)
But only **weak intermolecular bonds** between molecules
91
What are the characteristic properties of simple covalent molecules?
Low melting point
Usually insoluble in water
Do not conduct electricity (or heat)
92
Describe the structure/bonding in a giant covalent network
Contain billions of atoms with strong covalent bonds between them
93
What are the characteristic properties of giant covalent networks?
High melting point because all bonds in structure are strong covalent bonds
Insoluble in water
Do not conduct electricity (apart from graphite)
94
Describe the structure/bonding in a giant metallic lattice
All metals/metalic bonds have giant metallic lattice structure
Has a strong electrostatic attraction between the positive metal ions and the delocalised electrons between the ions
95
What are the characteristic properties of giant metallic lattices?
**High melting point** because there is **strong electrostatic attraction between ions + electrons**
**Insoluble in water**
**Conduct electricity when solid/molten** because **delocalised electrons are free to move** in response to voltage
96
What are delocalised electrons?
Electrons that are not associated with a particular atom
Instead are free to move over several atoms
97
In the giant ionic lattice/crystalline structure of NaCl, 6 Cl- atoms arrange themselves around 1 smaller Na+ atom
The Cl- atoms arrange themselves to be as far apart as possible. Suggest the name for the 3D arrangement they will take up
Octahedral
98
Group 1 metals have relatively low melting points
Use ideas about the charge and size of Group 1 ions to explain this
Group 1 ions have a small (+1) charge and have a relatively small ionic radius (compared with other ions in same period)
These 2 factors reduce the electrostatic attraction between the ions and delocalised electrons
They also only have 1 outer shell electron which can become delocalised, meaning electrostatic forces of attraction between ions and electrons are smaller
99
What is electron affinity?
Energy change when 1mol gaseous atoms aquires 1mol electrons from 1mol gasous anions
100
What is a complex ion?
Ion containing more than 1 atom
Charge is spread across whole ion
Contains covalent bonds
101
What is a precipitate?
A suspension of solid particles fromed by a chemical reaction in solution
102
What is a precipitation reaction?
Reaction between ions in solution that forms a precipitate
(suspension of solid particles/insoluble solid particles)
103
What does the prefix -ate tell you about the composition of an ion?
There are 1+ non-metal ions/atoms bonded to oxygen
e.g. sulfate -SO42-
104
What charge does the ClO ion have?
-1
105
What is the formula + charge of a phosphate ion?
PO43-
106
What charge does a lead ion have?
Pb2+
107
What charge does a nitrate ion have?
NO3-
108
What happens when ionic substances dissolve into solution/water?
The ions become surrounded by water + spread throughout the solution
They behave independantly of each other
109
Which ionic substances are soluble?
All compounds containing...
**Group 1 metals**
**Nitrate ions**
**Ammonium ions**
... are soluble
110
Which ionic substances are insoluble?
**Sulfates** of **Ba, Ca, Pb,** and **Ag**
**Halides** of **Ag** + Pb
**All carbonates *except*** those of **Group 1/ammonium ions**
**Hydroxides** containing **some Group 2**, **Al**, or **d-block ions**
111
Give the colour of precipitate (if one forms) that forms from the reaction of **OH-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
Ca2+ - White ppt
Ba2+ - White ppt
**Cu2+ - Pale blue ppt**
**Fe2+ - Green ppt**
**Fe3+ - Brown ppt**
Al3+ - White ppt
Pb2+ - White ppt
Zn2+ - White ppt
Ag+ - White ppt
112
Give the colour of precipitate (if one forms) that forms from the reaction of **SO42-**- ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
Ca2+ - White ppt
Ba2+ - White ppt
**Cu2+ - Soluble**
**Fe2+ - Soluble**
**Fe3+ - Soluble**
**Al3+ - Soluble**
Pb2+ - White ppt
**Zn2+ - Soluble**
Ag+ - White ppt
113
Give the colour of precipitate (if one forms) that forms from the reaction of **CO32-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Pb2+, Zn2+, and Ag+
Ca2+ - White ppt
Ba2+ - White ppt
**Cu2+ - Green ppt**
**Fe2+ - Green ppt**
Pb2+ - White ppt
Zn2+ - White ppt
Ag+ - White ppt
114
Give the colour of precipitate (if one forms) that forms from the reaction of **Cl-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
**Ca2+ - Soluble**
**Ba2+ - Soluble**
**Cu2+ - Soluble**
**Fe2+ - Soluble**
**Fe3+ - Soluble**
**Al3+ - Soluble**
Pb2+ - White ppt
**Zn2+ - Soluble**
Ag+ - White ppt
115
Give the colour of precipitate (if one forms) that forms from the reaction of **Cl-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
**Ca2+ - Soluble**
**Ba2+ - Soluble**
**Cu2+ - Soluble**
**Fe2+ - Soluble**
**Fe3+ - Soluble**
**Al3+ - Soluble**
Pb2+ - White ppt
**Zn2+ - Soluble**
Ag+ - White ppt
116
Give the colour of precipitate (if one forms) that forms from the reaction of **Cl-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
**Ca2+ - Soluble**
**Ba2+ - Soluble**
**Cu2+ - Soluble**
**Fe2+ - Soluble**
**Fe3+ - Soluble**
**Al3+ - Soluble**
Pb2+ - White ppt
**Zn2+ - Soluble**
Ag+ - White ppt
117
Give the colour of precipitate (if one forms) that forms from the reaction of **Br-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
Ca2+ - Soluble
Ba2+ - Soluble
Cu2+ - Soluble
Fe2+ - Soluble
Fe3+ - Soluble
Al3+ - Soluble
**Pb2+ - White ppt**
Zn2+ - Soluble
**Ag+ - Cream ppt**
118
Give the colour of precipitate (if one forms) that forms from the reaction of **I-** ions with:
Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+
Ca2+ - Soluble
Ba2+ - Soluble
Cu2+ - Soluble
Fe2+ - Soluble
Fe3+ - Soluble
Al3+ - Soluble
**Pb2+ - Yellow ppt**
Zn2+ - Soluble
**Ag+ - Yellow ppt**
119
The presence of which ions can be tested for by adding barium chloride solution?
(Solution containing Ba2+ ions)
Sulfate (SO42-) ions - white ppt formed
120
The presence of which ions can be tested for by adding silver nitrate solution (containing Ag+ ions)?
Shows presence of halide ions
White ppt forms for Cl-
Cream ppt forms for Br-
Yello ppt forms for I-
121
What are the 4 ways/reactions that can be used to make an ionic salt?
Acid + base/alkali → Salt + Water
Acid + Carbonate → Salt + Water + CO2
Acid + Metal → Salt + Hydrogen (MASH)
122
Give three examples of practical applications of precipitation reactions
Water treatment
Production of coloured pigments for paints/dyes
Identification of certain metal ions in solutions
123
What is relative isotopic mass?
The mass of one atom of an isotope compared to 1/12 of the mass of a 12C atom
(The same as Ar but for isotopes...)
124
What is empirical formula?
The simplest ratio of atoms in a compound
125
What is water of crystallisation?
Number of water molecules contained in an ionic lattice per molecule of salt
(i.e. how hydrates the salt is)
126
Hydrates cobalt (II) chloride has the formula CoCl2•xH2O
1.173g of hydrated cobalt chloride is heated to drive off the water of crytallisation. The mass remaining is 0.641g.
Calculate the formula of the hydrated salt
Calculate the **mass of water removed**:
1.173 - 0.641 = 0.532g
Calculate the **moles of water removed**:
0.532 / 18 = 0.02956mol
Calculate the **moles of anhydrous salt**:
0.641 / 129.9 = 0.00493mol
Calculate the **ratio**
* *moles water / moles anhydrous salt** =
0. 02956 / 0.00493 = 5.996 = 6
127
Calculate the percentage by mass of nitrogen in ammonium sulfate (NH4)2SO4
Calculate the **Mr of (NH4)2SO4** = 132.1
Calculate the **mass of N in 1 mol (NH4)2SO4**:
14 x 2 = 28
**% by mass = mass element in 1mol / Mr compound x 100**
(28/132.1) x 100 = 21.2%
128
What is the formula for percentage yield?
(Experimental yield / theoretical yield) x 100
**Yield** given by **moles** of reactants/products from **balanced equation**
129
Why might percentage yield be lower than expected?
**Loss of product** from reaction **vessels** (when transfering)
**Side reactions** (may create by-products)
**Impurities** in reactants
**Changes in temp + pressure** (may effect equilibrium)
**If the reaction is an equilibrium system**
130
What are the 3 main factors that affect 1st ionisation enthalpy?
**Atomic radii** - larger = lower
**Nuclear charge** - more protons = higher
**Electron shielding** - outer shells feel less electrostatic attraction to nucleus
131
How do group 2 metals react with water?
Group 2 metal + Water → Metal hydroxide + Hydrogen
M(s) + 2H2O(l) → M(OH)2(s) + H2(g)
132
How does the vigourousnes of Group 2 metals change as you go down the group?
The reactions become more vigorous as you go down the group
133
How do the hydroxides of Group 2 metals change as you go down the group?
Become increasing soluble and more alkaline
134
How do Group 2 metals react with oxygen (when heated)?
Metal + Oxygen → Metal oxide
2M(s) + O2(g) → 2MO(s)
135
What substances do Group 2 metal oxides react with?
What property does this give them?
They react with acids, so can act as bases
Metal oxide + Acid → Salt + Water
MO(s) + H2SO4(aq) → MSO4(aq) + H2O(l)
136
How do metal hydroxides react with acids?
Metal hydroxide + Acid → Salt + Water
M(OH)2(s/aq) + 2HCl(aq) → MCl2(aq) + 2H2O(l)
137
What happens to Group 2 metal carbonates when they are heated?
Give the general equation
Undergo **thermal decomposition**
Metal carbonate → Metal oxide + Carbon dioxide
MCO3(s) → MO(s) + CO2(g)
138
How does the thermal stability of Group 2 carbonates change going down the group?
As you go **down the group, thermal stability increases**
Means that **metals further down decompose at higher temperatures** than those further up the group
139
Why does the thermal stability of Group 2 carbonates increase as you go down the group?
**M2+ ions get larger as you go down** the group, so their **charge density is lower**
Because ions higher up the group have **greater charge densities, they polarise the carbonate ion more**
The **more polarised the carbonate ion, the more likely it is to break up** + form an oxide ion and CO2
140
What is a polarised ion?
A large (complex) ion that can have its electron distribution altered by small, highly-charged ions
This is known as polarisation
141
What is charge density?
The charge of an ion relative to its size
Mg2+ has a greater charge density than Ba2+ because, although they both have the same overall charge, Mg2+ is smaller
142
10.0cm3 of a 1.00moldm-3 solution of HCl was transfered to a volumetric flask and made up to 250cm3 with water.
What is the concentration of the diluted solution?
Calculate the **dilution factor**:
10/250 = 0.04
Use this to **calculate the new conc. of the diluted solution**:
0.04 x 1.00 = 0.0400moldm-3
143
What is an acid?
A substance that produces/donates H+ ions in a solution
144
What is a base?
A copound that reacts with an acid to produce water and a salt
Is a proton acceptor
145
What is an alkali?
A soluble base
Dissolves in water to produce hydroxide (OH-) ions
146
What is an oxonium ion?
What is its formula?
H3O+(aq)
Hydrogen ion bonded to a molecule of water
Created due to presence of H+ in acidic solutions (hence present in all acidic solutions)
147
Briefly describe how a soluble salt can be made
By reacting the appropriate acid and alkali together
The solid salt can then be produced by evaporating the excess solution/water
148
Breifly describe how an insoluble salt can be made
By a precipitation reaction
E.g. silver iodide can be made by reacting silver nutrate + potassium iodide
