Topic 456 Flashcards
Physical change
During a physical change, no new substances are produced. The chemical identity of the substances remains the same (they have the same chemical formulae before and after the process).
- Physical changes usually can be reversed easily
- E.g. Changes of state and formation of mixtures
Chemical change
During a chemical change new substances are formed and the chemical identities of the substances involved change but no atoms are created or destroyed. The combined mass of all the substances remains constant.
- Chemical changes are called chemical reactions
- Chemical changes usually are difficult or impossible to reverse
- E.g. Frying an egg, corrosion of copper roofs
Signs of a chemical change can include:
1. Flame/fire (burning or combustion, releasing heat or light energy)
2. Temperature change (increase or decrease)
3. Fizzing/ bubbles (effervescence) -> gas produced
4. Colour change (doesn’t include white to colourless because it is same in chemistry)
5. Precipitate formed (A solution turns into a suspension, opaque) -> an new insoluble solid is produced
6. New odour is detected
Chemical equations
Reactants -> products
State symbols indicate the physical states of the substances in a reaction:
(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous (dissolved in water)
Word equations use the names of the substances to describe a chemical reaction.
E.g. Hydrogen (g) + Oxygen (g) -> Water (l)
Whereas symbol equations use their chemical symbol.
E.g. H2 + O2 -> H2O
A balanced symbol equation shows the simplest ratio of the substances involved.
This means that the number of atoms of each element must be the same on both sides of the equation.
Oxygen transfer
Oxidation and reduction can be explained in terms of oxygen transfer and electron transfer.
Oxidation: Gain of oxygen by a substance in a chemical reaction
E.g. Mg + O -> MgO (Magnesium is oxidized)
Reduction: Loss of oxygen by a substance in a chemical reaction
E.g. 2 Al2O3 -> 4 Al +3 O2 (Aluminum oxide is reduced)
Oxidation & reduction take place in the same reaction. One reactant is oxidized, another is reduced. Oxygen is transferred. These reactions are called redox reactions.
Examples of redox reactions include combustion, corrosion of metals and extraction of metals from their oxides.
Electron transfer
Redox reactions can be explained more fully in terms of the transfer of electrons between two substances
Oxidation: Loss of electrons by a substance in a reaction
Reduction: Gain of electrons by a substance in a reaction
e.g. 2Mg + O2 → 2MgO
The equation is split into 2 half-equations
2Mg -> 2Mg 2+ +4e- (oxidation half-equation)
O2 + 4e- -> 2O 2- (reduction half-equation)
In reactions between metals and non-metals, metals are oxidized, and non-metals are reduced.
Displacement reacitons
Displacement reactions are redox reactions, where a more reactive element takes the place of a less reactive element in a compound.
In metal displacement, a more reactive metal takes the place of a less reactive metal.
GENERALLY: X (s) + YA (aq) -> Y (s) + XA (aq)
X is more reactive than Y
e.g Zinc (s) + copper(ii) sulfate (aq) → copper (s) + zinc sulfate (aq)
The sulfate ion takes no part in the reaction (spectator ions) and can be left out to make the ionic equation.
Zn + Cu2+ -> Zn2+ + Cu
The ionic equation can be split into 2 half-equations.
Zn -> Zn2+ + 2e- OX half equations
Cu2+ + 2e- -> Cu RED half equations
Oxidation numbers
-> a number assigned to each element in a compound showing how many electrons it has lost or gained
Pure elements have an oxidation number of zero.
For transition metal compounds the oxidation number of the metal is indicated by Roman numerals after the name of the metal. For metals it is always positive.
E.g. Copper (II) = Cu2+
During redox reactions, the oxidation number of some of the elements involved change.
An increase in the oxidation number of an element is oxidation.
A decrease in the oxidation number of an element is reduction.
e.g Iron(ii) chloride + chlorine → Iron(iii) chloride
The oxidation number of the iron has increased from +2 to +3 so Fe is oxidized.
Fe2+ → Fe3+ + e- loss of an electron.
EXOTHERMIC REACTIONS
Exothermic reaction is a reaction where energy is released from the substances into their surroundings. Chemical energy in the substances is converted to thermal energy in the surroundings. The products, therefore, have less energy than the reactants. Because the surroundings gain energy, the temperature increases.
The energy transferred during a reaction is called the enthalpy change ∆H. An exothermic reaction is a negative enthalpy change (since substances lose energy).
Exothermic: Freezing, condensation
ENDOTHERMIC REACTIONS
This is a reaction where energy is absorbed by the substances from their surroundings. Thermal energy (in the surroundings) is converted to chemical energy (in the substances). The products, therefore, have more energy than the reactants.
The energy absorbed can be in the form of heat, light or electrical energy. Because the surroundings lose energy, the temperature decreases.
An endothermic reaction is a positive enthalpy change (since substances gain energy).
Endothermic: melting, boiling
Energy bonding
Breaking bonds absorbs energy from the surroundings (Endothermic, ΔH is positive)
Forming bonds releases energy to the surroundings (exothermic, ΔH is negative)
If the total amount of energy absorbed to break bonds during a reaction is greater than the total amount released by bond making the reaction will be endothermic.
Every type of covalent bond has its own bond energy.
∆H = Total Energy of bonds broken - Total Energy of bonds made
If the reverse, the reaction will be exothermic
Activation energy
In most reactions, chemical bonds in the reactants must be broken before new bonds in the products are formed.
The breaking of these bonds requires energy. This is called the activation energy (Ea).
The activation energy is the minimum energy needed for bonds in the reactants to break. Many exothermic reactions still require an initial input of energy to start.
The activation energy can be shown on the reaction pathway diagram.
PEriodic table patterns
The Periodic Table arranges the elements in order of increasing atomic number.
The elements are organized into vertical columns called groups (I to VIII) and into horizontal rows called periods (1 to 7) according to their electronic structure.
Elements in the same Group have the same number of valence electron and this number is equal to the Group number (apart from Group VIII)
Elements in the same Period have the same number of occupied shells and this number is equal to the Period number. The outer shell gets progressively filled going across a Period.
On looking across each Period in turn, a periodic pattern is noticed, with the properties of the elements repeating themselves.
This is because elements in the same Group have similar chemical properties as they have the same electronic structure (same number of valence electrons).
Metals
Across a Period: Metallic properties decrease; Down a Group: Metallic properties increase. The metallic elements are located on the left and centre of the Periodic Table.
Most of the elements are metals, and they include all of Group I, Group II and the Transition elements.
All metals have the following physical properties (in isolation).
- Lustrous
- Electronically/Thermally conductive
- Malleable & Ductile
Most metals also have relatively high melting points and densities compared to most non-metals. Many metals are hard & strong.
Chemical Properties (behaves with others)
Metallic elements lose electrons to form positive cations.
Metals react with non-metals to form ionic compounds.
E.g. calcium (s) + oxygen (g) → calcium oxide (s)
Metal oxides have high melting points and are basic (react with acids).
Soluble metal oxides form alkaline solutions with water. These have a pH above 7 and turn Universal Indicator blue/purple.
Non-metals
The non-metallic elements are located on the right of the Periodic Table and include Group VIII, Group VII and most of Group VI.
Physical Properties:
- most do not conduct electricity (insulators) and are poor conductors of heat (except carbon: diamond and graphite)
- most have relatively low melting points & densities compared to metals
- many are gases at room temperature (VII elements)
- most solid non-metals are brittle & weak
Chemical Properties:
Non-metallic elements gain electrons to form negative anions.
Non-metals react with metals to form ionic compounds.
They also react with other non-metals to form covalent compounds.
sulfur (s) + oxygen (g) → sulfur dioxide (g)
Non-metal oxides usually have low melting points and are acidic.
Soluble non-metal oxides form acidic solutions with water. These have a pH below 7 and turn Universal Indicator red.
Alkali metals
-> The alkali metals are a group of soft metals with relatively low melting points and low densities. They are not alkali, but they produce hydrogen and an alkali when they react with water.
Going down the group, the melting point and density decreases, except potassium’s density.
Chemical properties:
- most reactive metals, more reactive going down the group
- one valence electron, easily lost to form 1+ ion
- corrode quickly when exposed to air (react with gases in the air) so are stored under oil
- react with non-metals to form ionic compounds; group I compounds are colourless and water-soluble
- produce hydrogen and an alkali when they react with water
(see notes for better understanding)
Alkali metals reacting with water
see notes
Halogens
-> typical non-metals that exists as diatomic molecules.
(see notes for colours!)
Chemical properties:
- reactive non-metals, decreases going down the group
- 7 valence electrons and easily gain one to form 1- ion
- Combine with metals to form ionic compounds known as halide salts
- Combine with other non-metals to form covalent compounds
Displacement reaction of halogens
See paper idk (see BOBO for help :))
Transition elements
Physical properties (compared to Group I metals):
- Have much higher melting points
- Are harder
- Are denser
(Zinc is not typical as its m.p. is much lower and is relatively soft)
Chemical properties:
- Less reactive than group I or II metals
- Have variable oxidation numbers (can form differently charged ions, indicated by roman numerals)
E.g. Iron (II) -> Fe2+ Iron (III) -> Fe3+
(Zinc can only form 2+ ions tho)
- Transition metal compounds and their solutions are usually coloured
Cu2+ -> blue or blue green Fe3+ -> orange brown
Fe2+ -> pale green (you need to know this!)
(Zinc compounds are colourless)
- The elements and their compounds are good catalysts
Noble gases
-> Colourless, monoatomic gases at room temperature
Going down the group, the boiling point and density increases.
All noble gases are found in less than 1% of the atmosphere combined.
They are inert (completely unreactive) as they have a full valence shell and so are completely stable. They do not form any chemical compounds.
They can be used to create an inert atmosphere:
- In light bulbs for lamps
- In extraction of certain metals
(other gases would react)
