Paper 1 Flashcards
1
Q
Why do isotopes have the same chemical properties?
A
This is because they contain the same number of electrons. Electrons determine chemical properties of atoms.
2
Q
John Dalton
A
John Dalton described atoms as tiny spheres that could not be divided
3
Q
JJ Thomspson
A
- J.J. Thomson discovered the plum pudding model
- It is a ball of positive charge with electrons embedded in the ball of positive charge.
- No empty space in the atom
- Mass spread throughout
4
Q
Ernest Rutherford:
A
- Ernest Rutherford’s alpha scattering experiment led to the plum pudding model being replaced by nuclear model.
- Positive charge concentrated at the centre
- Most of the atom is empty space
- Mass is concentrated at the centre
5
Q
Niels Bohr
A
- Niels Bohr adapted the nuclear model by suggesting that the electrons orbit the nucleus at specific distances.
- The electrons are on electron shells/energy level.
6
Q
Later experiments:
A
- Later experiments led to the idea that the positive charge of any nucleus could be subdivided into a whole number of smaller particles, each particle having the same amount of positive charge.
- The name proton was given to these particles.
- James Chadwick provided the evidence for the existence of neutrons in the nucleus.
- This was about 20 years after the nucleus became an accepted scientific idea.
- This could help explain the existence of isotopes.
7
Q
Modern periodic table
A
- Arranged in order of atomic number, elements with similar properties are in columns, known as groups.
- Periodic table- siimilar properties occur at regular intervals.
- Elements in same group - same number of electrons in their outer shell- gives them similar chemical properties.
- Elements in the same period have the same number of electron shells
- As you go down the group, the number of shells increases.
- Metals are found on the left-hand side of the periodic table whilst non-metals on the right-hand side
- Most of the elements found on the period table are metals
8
Q
Properties of metals:
A
- High density
- High melting point
- Shiny
- Malleable
- Conducts electricity and heat
9
Q
Properties of non-metals:
A
- Low density
- Low melting point
- Dull
- Brittle
- Does not conduct electricity nor heat
10
Q
What were early periodic tables arranged in order of?
A
In order of atomic weight.
11
Q
Newland’s periodic table:
A
- Have more than one element in a box
- Missing Noble gases
- Arranged elements in order of atomic weight
DISADVANTAGES:
* He did not leave gaps for undiscovered elements
* He had many dissimilar elements in a column
12
Q
Mendeleev’s periodic table:
A
- Have more than one element in a box
- Missing Noble gases
- Arranged elements in order of atomic weight
Advantages of Mendeleev
* He left gaps for undiscovered elements
* He changed the order of some elements
13
Q
Why was Mendeleev’s periodic table accepted?
A
- He predicted properties of missing elements
- Elements with properties predicted by Mendeleev were discovered
- Elements were discovered which fitted the gaps
14
Q
Group 0 (Nole gases)
A
- Noble gases exist as single atoms (monoatomic) as they have a full outer shell, which means they do not lose or gain electrons.
- They are unreactive because they have a full outer shell and they do not need to lose
or gain electrons. - The boiling points of the noble gases increase with increasing relative atomic mass (going down the group).
15
Q
Group 1 metals:
A
- They react with non-metals to form ionic compounds
- The compounds are white solids which dissolve in water to form colourless solutions.
- The alkali metals all have low densities.
- The first three alkali metals (Li, Na & K) are all less dense than water – this is why they float on the surface.
16
Q
Observation of the first three alkali metals with water
A
- The alkali metals react with water to form a metal hydroxide and hydrogen gas.
- When universal indicator is added to the metal hydroxide solution, it turns purple as it contains OH-(hydroxide) ions.
17
Q
Lithium reaction with water:
A
- Fizzing/ bubbles
- Moves on the surface of the water
- Gets smaller in size
18
Q
Sodium reaction with water:
A
- Fizzing/ bubbles
- Moves quickly on the surface of water
- Gets smaller in size
19
Q
Potassium reaction with water:
A
- Fizzing/ bubbles
- Moves very quickly on the surface of water
- Gets smaller in size
- Lilac flame
20
Q
Why does the reactivity of group 1 metals increases down the group?
A
- More electron shells down the group
- Weaker attraction between the nucleus and the outer shell electron
- Easy to lose the outer electron
21
Q
Halogens:
A
- All have 7 electrons on their outer shell
- They have coloured vapours
- Consist of molecules which are made up of pairs of atoms (diatomic) For example, Cl2, F2, Br2
22
Q
Why does the melting point and boiling point increase down group 7?
A
- The melting and boiling point increases as you go down group 7 because the relative molecular mass increases.
23
Q
Why does reactivity decrease down group 7?
A
- More electron shells
- Weaker attraction between the nucleus and the outer shell electron
- Harder to gain an electron
24
Q
Halogen- displacement reaction:
A
A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.
25
Compared with group 1 elements, the transition metals:
* Have higher melting points
* Have higher densities
* Are stronger
* Are harder
* Are muss less reactive- so do not react as vigorously with water and oxygen
26
Properties of transition metals:
* Have **ions** with **different charges**
* Form **coloured compounds**
* Are useful as **catalysts**
27
Atoms:
* **All substances** are made of **atoms.**
* An atom is the **smallest part of an element** that can **exist.**
28
Element:
A substance made form **one type of atom**
29
Compounds:
* Contains **atoms** of **two or more different elements** **chemically bonded together.**
* Is **difficult** to separate.
30
Mixture:
* A **mixture** is when **two or more different elements** or **compounds** that are **not** chemically bonded together.
31
Key definitions- seperation techniques:
* **Soluble –** it dissolves.
* **Insoluble –** it will not dissolve. Examples of insoluble substances include mud and sand.
* **Solute –** a solid being dissolved. Example include salt and sugar.
* **Solvent –** a liquid a solute can be dissolved into.
* **Solution –** a mixture of a solute and solvent that does not separate.
32
Filtration:
Used to seperate an insoluble solid from a liquid
33
Simple distillation:
* Used to seperate liquids with very diffferent boiling points
* Used to seperate a liquid from a soluble solid
34
Fractional distillation:
Used to seperate solutions with similar boiling points
35
Chromatography:
Chromatography is used to **separate mixtures** of **coloured compounds.**
36
Chromatography practical:
1. Use a **pencil** and **ruler** to draw **startline** and **solvent** front on a **chromatography paper**
2. Fill a **beaker** with the **solvent** and ensure that the **startline** is **above** the **solvent.**
3. Place **dots** of the **samples** to be tested on the **startline**
4. Place the **chromatography paper** in the **beaker** and add **lid.**
5. Allow the **solvent** to move **through** the paper
6. Take out the **paper** and allow it to **dry.**
7. Analyse the separation.
37
Two phases of chromatography:
1. A stationary phase
2. A mobile phase
* The **paper** is the **stationary** phase
* The **solvent** is the **mobile** phase
38
Chromatography key points:
* **Start line** drawn in **pencil** to prevent it from **dissolving in the solvent**
* The **spots** should be **above the solvent** to prevent it from **dissolving in the solvent**
* A lid is used to prevent solvent from evaporating
* The **most soluble colour** will travel the **furthest** and the **least soluble** will travel the **least.**
* If a **colour is insoluble,** it will **not move** therefore the **solvent** has to be **changed.**
* A **pure sample** only has **one dot in a vertical line.**
* The **sample** is a **mixture** if there is **more than one dot in a vertical line.**
39
How the different dyes separated by paper chromatography? 4 marks
* Solvent moves through paper
* Different dyes have different solubilities in solvent
* Different dyes have different attractions for the paper
* So are carried different distances.
40
Rf value:
Distance moveed by substance/ distance moved by solvent
41
Structure of ionic compounds:
* A giant ionic lattice
42
Ionic lattice:
* The **strong ionic bonds** require **lots of energy** to **break**
* Ionic compounds are **mostly white solids.** Just like salt (NaCl).
* Ionic compounds **easily dissolve in water** to form **aqueous solutions.**
* As an **ionic compound dissolves,** its ionic **lattice breaks up**, and the ions are **free to move.**
43
**Why** can **solid ionic compounds** not conduct **electricity?**
- Ions are **fixed** in position
- Ions can only **vibrate** (they can’t move)
- Ions can’t carry **electrical charge**
44
Why can **molten/ aqeous ionic** compounds conduct electricity?
* Ions are **free** to **move**
* Ions can carry an **electrical charge**
45
Explain why sodium chloride has a high mp?
* Giant ionic lattice
* SEABOCI
* A lot of energy is needed to overcome the forces
46
What is a compound ion?
Groups of elements that have a charge
47
Small covalent molecule:
* Small covalent molecules are **gases** at **room temperature**
* **Weak intermolecular forces** between **molecules** which require **little energy to break.**
48
Why is **ammonia** a **gas** at room temperature?
* Ammonia is a **small covalent molecule**
* Ammonia is **made up of small molecules**
* There are **weak intermolecular forces** beetween bonds
* **Little energy** is needed to **break bonds**
49
Graphite (giant covalent structure):
* A giant covalent structure
* Only made from carbon atoms
* Arranged in hexagonal layers
* There is weak intermolecular forces between layers
* Each carbon atom is bonded to three other carbon atoms by STRONG covalent bonds
50
Uses of graphite:
* Lubricants
* In pencils (lead)
* Electrodes
51
Why is graphite **soft?**
* There are only **weak intermolecular forces** between **layers.**
* The **layers** (of carbon atoms) in graphite can **slide.**
52
Why does graphite **conduct electricity?**
* Each **carbon atom** is bonded to **three other carbon atoms.**
* **One delocalised electron** per **carbon atom** which **moves** throughout the **structure** and **carry electrical charge.**
53
Why does **graphite** conduct **thermal energy?**
* Each **carbon** is bonded to **three other carbon atoms. **
* **One delocalised electron** per **carbon atom** which **moves** throughout the structure and can **transfer thermal energy.**
54
Why does **graphite** have **high melting point?**
* **Giant** covalent structure
* **Each carbon atom** is bonded to **three other** carbon atoms by **strong covalent bonds.**
* Need a **lot of energy** to break the **bonds.**
55
Diamond:
**Structure and bonding**
* A **giant** covalent structure.
* Made only from **carbon atoms.**
* **Each carbon atom** is bonded to **four** other carbon atoms.
56
Uses of diamond:
* Laser beam
* Cutting tools and drills
* Jewellery
57
Silicon dioxide:
**Structure and bonding**
* **Giant** covalent structure.
* Made from **silicon** and **oxygen** atoms.
* Each **oxy** atom- bonded to **2 silicon** atoms by **strong covalent bonds.**
* Each **silicon** atom- bonded to **four** other **oxygen** atoms by **strong covalent bonds.**
* Has a formula **SiO2.**
58
Uses of silicon dioxide:
Lining for furnaces- due to high mp.
59
Graphene:
* Giant covalent structure
* It is made of a single layer of graphite.
* Each carbon atom forms three strong covalent bonds.
* It has a low density
* Useful in elctronic devices- it is transparent
60
Fullerenes:
**HOLLOW SHAPED MOLECULES** THAT ARE ONLY MADE FROM **CARBON ATOMS**
* Drug delivery
* Lubricants- spherical, so can roll
* Catalysts- large SA to V ratio
61
Carbon nanotube:
* **Cylindrical fullerenes** made from **ONLY carbon atoms.**
* They are **strong** and have a **low density.**
* **Each** carbon atom in carbon nanotubes forms **three strong covalent bonds.**
* This means there are **delocalised electrons**- carry an **electrical charge**
* They also have a **high tensile strength** (don’t break when stretched).
* **Nanotechnology**
* **Electronics**
* **Materials** as it has a **low density.**
62
Metals:
* Giant structure of positively charged ions in a lattice.
* Delocalised electrons move throughout the whole structure.
* Strong electrostatic attraction between positive ions and negative electrons.
63
Why are metals malleable?
Layers of atoms can slide
64
Why do **metals** have high mp and bp?
* **Giant lattice.**
* Strong electrostatic force of attraction between **positive ion** and **negative electrons.**
* A **lot of energy** is required to **break the force.**
65
Nanoparticles:
**Very high SA to V ratio**
**USES:**
* Huge surface area to volume ratio, so they could make new **catalysts**
* **Drug delivery**
* Some nanoparticles conduct electricity, so they can be used in **tiny electric circuits for computer chips**
* They are also used in **cosmetics.** Used to improve **moisturisers** without making them
**really oily.**
66
Risk of nanoparticles:
* Some nanoparticles could be toxic
* Cell damage to the body
* Un-discovered harmful side effects
67
Coarse particles diameter:
2.5 * 10 ^-6 m
68
Fine particles diameter:
1 * 10^-7 m
69
Nanoparticles diameter:
1 * 10^-9 m.
70
Law of conservation of mass:
* **No atoms** are **lost** or **made** during **chemical reaction**
* **Mass** of the **products** equals the mass of **reactants**
71
Thermal decomposition of copper carbonate:
Copper carbonate → Copper oxide + carbon dioxide
72
When is atom economy 100%?
When only one product is formed
73
Atom economy:
(Total mass of desired product atoms/ total mass of reactant atoms) * 100
74
Mass of a sinle atom in grams:
Mass of a single atom in g = Relative atomic mass / (6.02 × 1023)
75
Percentage yield:
(Mass of product actually made/ maximum theoretical mass) * 100
76
Moles of GAS:
Volume of gas dm^3 / 24 dm^3
77
Cells:
* A cell is a system that contains chemicals which react to produce electricity.
* A simple cell can be made by connecting two different metals in contact with an electrolyte.
78
Alkaline batteries:
* Alkaline batteries contain **hydroxide ions**
* They are used in **remote controls**
* Over time, the **ions** in the **electrolyte** and the **metal ions** on the electrode gets **used up and turns into products**
* Once any one of the reactants get used up, the **reaction can’t happen** and so **no electricity is produced**
* Reactions in an alkaline battery is **irreversible** therefore they are **non-rechargeable.**
79
Hydrogen fuel cells:
Hydrogen + oxygen = water + energy
80
**Anode-** hydrogen fuel cell:
OXIDATION:
**H2→2H+ + 2e-**
81
**Cathode-** hydroge fuel cells:
**REDUCTION:**
2H2 + O2 → 2H2O
82
Hydrogen fuel celss + and -:
**Advantages**
* No pollutants-Produces only water
* Hydrogen can be made renewable if made by electrolysis using renewable energy.
**Disadvantages**
* Hydrogen is highly flammable
* Hydrogen not renewable if produced using fossil fuels
* Not many hydrogen filling stations
83
Rechargeable batteries + and -:
**Advantages:**
* Charging points more widely available
* Rechargable
**Disadvantages:**
* May release toxic chemicals on disposal
* Rechargeable batteries have finite life-time
* Can catch fire
84
Bond energies:
* **Energy supplied** to **break** chemical bonds so **bond breaking** is an **endothermic process**
* **Making new bonds** transfers **energy** to the **surrounding** therefore is an **exothermic process.**
85
Endothermic reaction:
**Energy** needed to **break existing bonds** is **greater** than the **energy released from forming new bonds.**
86
Exothermic reaction:
**Energy released** from **forming new bonds** is **greater** than the energy **needed** to **break existing bonds.**
87
Metals and extraction:
* Metals that are **less reactive** than carbon can be **extracted** from their **oxide** by **reduction with carbon.**
* Metals **more reactive** than carbon are **extracted by electrolysis.**
