CHEMICAL OR PHYSICAL CHANGE
Physical change- state change of solid to liquid, easily reversed, chemistry of substance is unchanged.
Chemical change- clues: change in colour/heat taken in or given out/bubbles form, difficult to reverse. Changes the chemical bonds between atoms allowing the arrangement of atoms to make new substances.
Examples:
Copper Sulphate: Before=blue crystals. After heated=becomes solid white. The colour changed, which indicates a chemical reaction.
Sodium Chloride: Before=white crystals. After heated=becomes powdery. This is a physical change, as the crystals only became powdery.
THE PARTICLE NATURE OF MATTER
Matter is defined as anything that has mass and takes up space. There are 3 physical states: Solid, Liquid and Gas.
The Solid State:
Particles are packed tightly together. Particles vibrate but are not free to move.
Kinetic theory (it assumes all particles are spheres which is not true, but easier to explain theory):
In solids, particles move a little or vibrate about fixed positions.
They have a regular structure. This explains why many solids are crystals. The same substance will have the same crystal structure.
Have definite volume and shapes which can be affected by changes in temperature (expansion/contraction).
Attractive intermolecular forces hold the particles together.
(See figure 1 for diagram of how particles moves/arranged in a solid)
The Liquid State:
Particles move more freely
Constantly moving. Particles are close together and undergo frequent collisions with each other, but intermolecular forces still operate and are weaker than in solids) Particles have more energy than in solids
There is a disordered arrangement They take the shape of the container They are less dense than solids
They cannot be compressed (or if so, VERY slightly)
(See figure 2 for diagram of how particles moves/arranged in a liquid)
The Gas State:
There are no forces binding particles together
The move independently, rapidly and randomly-intermolecular forces and particle size can be ignored.
Their density is very low compared to solids and liquids.
Rapid movement allow the gas to diffuse to areas of low concentration.
They have no shape but fill the entire volume of the container. (volume of gas expands when you heat it)
(See figure3 for diagram of how particles moves/arranged in a gas)
(See figure 4 for diagram of the overview of matter) ENERGY CHANGES TO DO WITH CHANGE OF STATE
(See figure 5 for state change diagram)
MELTING
Solids start heating up, and gain enough energy to push past each other.
This causes the solid to expand in volume (except water-ice is actually frozen water with air molecules in it, that‟s why it can float on liquid water)
In the end the forces of attraction are weakened enough for the solid to turn to liquid. The normal structure becomes lost.
This is the melting point and the temperature will no longer rise until all the solid has melted.
The higher the melting point, the stronger the forces of attraction between particles. EVAPORATION
As the liquid gains more energy some of the particles on the surface gain enough energy to escape the attractive forces and leave, the liquid becomes gas.
When bubbles start to form within the liquid as the gas forms so quickly, you have reached boiling point.
This is a physical change
At boiling point the pressure that the gas is at is atmospheric CONDENSATION/FREEZING
When we condense and freeze attractive forces re establishes themselves and energy is given out.
These are physical changes to a substance as no new substance is formed. SUBLIMATION
A few solids such as carbon dioxide (dry ice) do not melt when they are heated. Instead they turn directly into a gas when they are heated. This is called sublimation.
This happens at a specific temperature. Iodine also sublimes.
ENERGY GIVEN OUT Gas-Liquid-Solid
ENERGY GAINED
OTHER CHANGE IN PHYSICAL CONDITIONS THAT CHANGE A STATE WITH SAME TEMPERATURE Changing the pressure:
State change can occur by changing the pressure. E.g. hairspray aerosol.
Also, pressure (not just heat) can cause ice to melt. Pressure can cause a change in state. I.e. ice melts when pressure is applied-skating.
On the top of high mountains, boiling point changes as air pressure changes.
Consists of only 1 substance. It has no contaminating properties. A pure substance melts and boils at a definite temperature.
These are predictable and precise and can be used to check for purity. This is extensively used to test for the purity of drugs both legal and illegal! EFFECTS OF IMPURITIES
Presence of impurities in a substance lowers the melting point and Raises the boiling point of the substance
An impurity reduces the „sharpness‟ of the melting point and boiling point, so state changes over a range of temperatures (a mixture of all the different substances‟ state change points) Note: Salt is put on roads in winter because it makes the liquid water on the roads be able to stay a liquid at lower temperatures because the melting/freezing point is lower.
LATENT HEAT
Latent heat is the energy required or released when a substance changes state WITHOUT a change in temperature.
(See figure 6 for a diagram on energy changes during heating and cooling)
For wax, the heating curve looks like (See diagram 7). This is because wax is an impure substance, so it doesn‟t have a definite or specific melting point. It melts over a range of temperatures.
COOLING STEARIC ACID CONCLUSION
Energy (in the form of heat) is given out as attractive forces between particles in a liquid are set up to become a solid. This keeps the „freezing point‟ artificially high for a time.
Breaking attractive forces TAKES IN energy. Making attractive forces RELEASES energy. DIFFUSION IN LIQUIDS AND GASES BUT NOT IN SOLIDS
Diffusion involves the movement of particles from a region of higher concentration towards a region of lower concentration. Eventually the particles are spread evenly-the concentration is the same throughout.
Diffusion does not take place in solids
Diffusion in liquids is much slower than in gases
Remember, in liquids and gases, particles will spread out to fill the space available.
This because particles are always moving, which why there is movement between high and low concentrations or particles.
Solids cannot diffuse because the particles in the solids cannot move, however liquid particles can move,
Gas diffuses faster than liquid particles because gas can move more because attractive bonds are weaker.
INVESTIGATING DIFFUSION
(See diagram 8 for lead nitrate + potassium iodide reaction) What is happening?
The crystals dissolve in the water and diffuse from an area of high concentration to an area of low concentration. This can occur because particles in a liquid can move. The point at which the crystals meet forms a yellow line of PbI2. This line is closer to the lead crystals than to the potassium crystal because lead is more dense and cannot move as far (having more molecular mass).
RATES OF REACTION
What factors lead to a successful reaction? Reactions can be fast or slow e.g. Fast=fireworks or Slow=apple turning brown.
Rates of reaction-it is concerned with the dynamics of chemical reactions such as the way reactions take place and the rate (speed) of the process.
It is vital in industrial chemistry because the time and energy required for a reaction are of great economic importance.
The basis of the study is COLLISION THEORY COLLISION THEORY
There are a few things that must happen for a reaction to take place.
Collision theory states that: Particles must COLLIDE before a reaction can take place. Not all collisions lead to a reaction
Reactants must possess at least a minimum amount of energy ACTIVATION ENERGY (units kJmol(-1))
Particles must approach each other with the correct orientation STEARIC EFFECT (head on collisions are ideal) (ORIENTATION)
If they move too slowly (don‟t possess enough energy to over the activation energy barrier) or at the wrong angle-they bounce off each other and don‟t react)
According to collision theory, to increase the rate of reaction you therefore need… More frequent collisions-increase particle speed or have more particles present
More successful collisions-give particles more energy OR lower the activation energy (the amount of energy they need to react, to break the bonds)
NOTE: Only a miniscule percentage of collisions lead to a reaction, otherwise the atmosphere would consist of nitrogen oxide!)
MEXO BENDO
For a chemical reaction to occur bonds between particles first have to be broken. Energy is taken in. This an ENDOTHERMIC PROCESS
„Bendo‟ is breaking in endothermic
When products are formed in a chemical reaction, bonds have to be made. Energy is given out. This is an EXOTHERMIC PROCESS
„Mexo‟ is making is exothermic.
INCREASING THE RATE OF CHEMICAL REACTIONS (6 ways) Increase the temperature
Increase the concentration of reactants Increase the pressure of any gases Increase the surface area of any solids Add a catalyst
Shine light
TEMPERATURE AND REACTION RATE
Increasing temperature increases the rate of reaction
Particles get more energy so they can overcome the energy barrier, THE ACTIVATION ENERGY
Particles speed also increase the number of SUCESSFUL collisions in a given time (the FREQUENCY of the collision)
More energymore collisionsmore successful collisions (See figure 10 for a diagram of distribution energy curves)
Only particles with greater or equal energy to the Activation energy can react when collisions occurs.
Temperature T2>T1
Increasing the temperature alters the distribution Get a shift to higher energies/velocities
Curve gets broader and flatter due to the greater spread of values
Area under the curve stays constant-it corresponds to the total number of particles Increasing temperature gives more particles an energy greater than Ea, more
reactants are able to overcome the energy barrier and form products. A small rise in temperature can lead to a large increase in rate.
EFFECT OF SURFACE AREA ON REACTION RATE
The bigger the surface area, the more particles are exposed and so can react. (See figure 11 for diagrams of increased surface area effect)
Increasing surface area increases the amount of particles that are in contact with each other. Therefore MORE collisions occur and there is an INCCREASES FREQUENCY of SUCCESSFUL COLLISIONS and a faster reaction.
(See figure 12 for a graph showing the rate of reaction with changes in surface area)
EFFECT OF CONCENTRAION ON THE RATE OF A REACTION
Increasing concentration of a reactant leads to more frequent collisions, which leads to an increased rate of reaction.
(See figure 13 for a diagram of concentration effect) Exam speak Summary-increasing the concentration
This increase the number of particles in the SAME VOLUME. Therefore MORE collisions occur and there is an INCREASED FREQUENCY of SUCESSSFUL COLLISIONS and a faster reaction.
INCREASING PRESSURE
(For gases, increasing the pressure has the same effect as increasing the concentration) Increasing the pressure forces gas particles closer together, which increase the chance of successful collisions, so the rate increases
The more particles that are in a given volume, the greater the pressure, so the more frequent collisions and the greater the chance of a reaction.
RATE CHANGE DURING A REACTION
(Reaction rate: how many successful collisions in a certain amount of time)
Reactions are fastest at the start and get slower as the concentration of reactants drops (See figure 14 for a diagram of a graph of the below information)
In a reaction such as A+2BC the concentration might change as shown: Reactants (A and B)-Concentration decreases with time
Product C-Concentration increases with time
The steeper the curve the faster the rate of reaction
Reactions start off quickly because of the greater likelihood of collisions Reactions slow down with time as there are fewer reactants to collide
RATE: How much concentration changes with time. It is the equivalent of velocity. (See figure 15 for a diagram of gradient vs. reaction rate)
The rate of change of concentration is found from the slope (gradient) of the curve The slope at the start of the reaction will give the initial rate
The slope gets less (rate is slowing down) as the reaction proceeds MEASURING THE RATE
EXPERIMENTAL INVESTIGATIONS
Measuring the volume of gas evolved
Measuring the change in mass if a gas is evolved Using a colorimeter or UV/visible spectrophotometer
Volume of gas can be measured to monitor the rate of a reaction. Volume against time graph.
Calcium carbonate (marble chips) +Hydrochloric acidCalcium chloride + water + carbon dioxide CaCO3(s) + 2Hcl(aq)CaCl2(aq)+H2O)l)+CO2(g)
The gas given off can be collected in a syringe and readings taken every 30 seconds or so (See figure 16 for a diagram of collecting gas apparatus)
Mass lost during a reaction can be used to monitor the rate Mass loss against time graph
Magnesium+ Hydrochloric acidMagnesium chloride +Hydrogen Mg(s)+2Hcl(aq)(MgCl2(aq)+H2(g)
As the gas given off leaves the flask the total mass of the fluid decreases slightly. Readings of the mass (g) can be taken. Typically at 1 min intervals.
(See figure 17 for a diagram of measuring mass loss apparatus)
Colorimetric measurements can be done by the eye or with a suitable meter (Light meter) Concentration or temperature against time graph
Sodium thiosulphate + Hydrochloric acidSodium Chloride + water+ Sulphur dioxide+ Solid sulphur Na2S2O3(aq)+2Hcl(aq)2NaCl(aq)+H2O(l)+SO2(g)+S(s)
The effect of changing conditions such as temperature of concentration can be studied by measuring how long it takes to produce enough sulphur to make the solution opaque (See figure 18 for a diagram of colorimetric apparatus)
NOTE: (See figure 19 for a handy starter graph on reaction rates compared) Q. Which graph would you expect to see for a rate of reaction experiment? A. Reaction was fast at the start then slowed down.
Q How would you calculate the rate of reaction?
A. You would find the gradient of a certain part of the graph. The gradient tells us how fast the reaction is by showing us how much gas (mL) is produced per second. y/x
INSERT PAGE OF CUT OUTS ABOUT RATE OF REACTION-FACTS-REASONS-MODELS CATALYST AFFECTS THE RATE OF REACTION
Adding a catalyst provides an alternative lower energy pathway. Therefore MORE particles have SUFFIECIENT energy to overcome the lower activation barrier. Therefore there is an increased FREQUENCY of SUCCESSFUL COLLISOONS and a faster reaction.
Catalysts remain unchanged at the end of a reaction and can be reused. (See figure 20 for a graph of how catalysts affect reaction rate)
ELEMENTS COMPOUNDS AND MIXTURES
(L.O. Describe the difference between elements, compounds and mixtures. Describe the difference between metals and non-metals.)
Elements are materials made up of one type of atom only.
Compounds are produced when elements combine during a chemical reaction It follows that compounds contain two or more types of atom
(See figure 21 for a simple diagram of element then compound) Compounds are produced when elements combine
The atoms in compounds are not mixed together
They become bonded together during a chemical reaction
Because of this: compounds have properties that are very different to the elements that they are made from. (See figure 22 for a diagram of this)
Mixtures are, as the name indicates, mixed rather than reacted together. This means unlike compounds:
They do not have any particular proportions of the various ingredients
Their properties are often an „average‟ of the properties of the ingredients (e.g. a mixture of a black and white powder is grey!)
They are mixed, not bonded, and so are usually not too hard to separate back into their ingredients (for example, it is easy to get salt back from sea water)
There are no chemical bonds in mixtures SYMBOLS
Chemists have agreed symbols that they use to represent elements
Many of the common elements just use the first letter of the name E.g. F=Fluorine Others use two letters E.g. Mg=Magnesium Ca=Calcium
Others are surprising E.g. Pb=Lead Ag=Silver Fe=Iron
If the symbol has just one letter make it a capital letter E.g. Nitrogen is N not n
If the symbol has two letters make the first a capital letter and the second lowercase E.g. Cobalt=Co not CO (carbon monoxide)
METALS VS NON METALS
Metal Non Metal
Strong
Malleable (easily shaped) High densities
Can be magnetic (only iron, cobalt, nickel) Conduct heat/electricity
Ductile (can be drawn out to wires)
Sonorous (makes a good sound when you hit it)
Dull
Brittle (Like glass) Low densities
Bromine is the liquid non metal, Mercury is a liquid metal (at room temperature) Bromine element exists as a molecule (two or more atoms joined together) ALLOYS
Metals can be alloyed to make them more useful.
Different metals are used for different properties e.g. A bell needs a shiny, sonorous and malleable metal, A saucepan needs the properties shiny (shiny surfaces reflect heat), high melting point, conductor, malleable)
(See figure 23 for a diagram of an alloy versus pure metal) What is an alloy?
Imagine you pushed one layer of atoms in a metal. What would happen? The layer will slide over the others-metals are malleable because atoms easily slide over each other.
Imagine you pushed one layer of atoms in an alloy. What would
happen? The layers won‟t slide over the others-alloys are harder and do not change shape easily. Generally alloys are harder, because the atoms no longer slide over easily.
Note: In carbon steels, a tiny amount of carbon (around 0.03%-1.5% is mixed with iron. If you add too much, you ruin the metal structure and it goes too brittle.
Rate of reaction: How fast reactants are changed into products
Pure iron and pure copper is too soft to be useful. So copper can be allowed with tin to make bronze, and with zinc to make brass.
ATOMIC STRUCTURE
Atoms are very small (if a helium atom is as big as a full stop, then a mouse is as big as the earth)
Subatomic particles: protons, neutrons, electrons
Most of the atom is just empty space (an atom can thought of as a foot stadium. The nucleus is the ball in the centre, and the electrons would be flies flying around. The rest is empty)
Subatomic particles in more detail
Subatomic Particle Relative Charge Relative Mass Common Depiction Proton
(found in the nucleus)
+1 1amu (atomic mass
unit) Neutron
(found in the nucleus) 0 1 amu
Electron (found in the electron shells around the nucleus) -1 1x10^(-5) (1/1840) amu ALLOYS ALLOYS Carbon Steels Copper Alloys Alloy Steels
LOW CARBON STEEL
0.5% of less carbon called mild steel Cheap, soft, quite strong, not brittle and easily shaped
Used in construction, bridges and large structures
HIGH CARBON STEEL More than 0.5% carbon Very strong/tough but brittle Used to make razor blades and cutting tools
HIGH ALLOY STEEL Most expensive
Contain higher amounts of other metals (e.g. chromium steel has 12-15% chromium, known as stainless steel
Tough, does not corrode
Used for cutlery and cooking utensils
BRASS
Contains 35% zinc and 65% copper Stronger and lower melting point than either metal
Strong, can be easily shaped Used got jewellery,instruments (good sonorous ability), electrical connectors
BRONZE
Contains 10% tin and 90% copper
Strong, does not corrode easily Used for structures, springs and machine parts
LOW ALLOY STEEL
More expensive than carbon steels Contain 1-5% other metals (e.g. chromium, manganese, nickel, titanium)
Tough, Hard
Used for edges of high speed cutting tools
Lithium
Mass number=Nucleon number (the big number is always the mass number, sometimes they have the numbers upside down)
Atomic number=proton number (number of protons. Number of protons=number of electrons)
Lithium has 3 electrons (+3 charge), 3 protons (-3 charge), and 4 protons (no effect) Number of protons=number of electrons. This is because the atom is neutral. The charge balances out.
Atoms can gain and lose electrons (they become ions). This changes the overall charge on the atom.
Number of protons „defines‟ an element (the number of protons NEVER changes) Atomic number does NOT always equal the number of neutrons. This can change even in atoms of the same element. These are called isotopes.
(See figure 24 for a diagram of some isotopes) The smallest whole particles are ATOMS
An atom consists of a central nucleus surrounded by much smaller electrons ELECTRONIC STRUCTURE AND THE PERIODIC TABLE
The horizontal rows are called the periods
Elements in the same period all have the same number of electron shells The vertical columns are called groups
Elements in the same group all have the same number of electrons in their outer shell so they all have similar properties
For GCSE, you only need to know how the rules apply for the first 20 elements (HCa) Electrons are arranged in shells or orbital‟s
The 1st electron shell contains a maximum of 2 electrons The 2nd electron shell contains a maximum of 8 electrons The 3rd electron shell contains a maximum of 8 electrons
SHELL 1 SHELL 2 SHELL 3 SHELL 4
Hydrogen 1 Helium 2 Lithium 2 1 Berylium 2 2 Boron 2 3 Carbon 2 4 Nitrogen 2 5 Oxygen 2 6 Fluorine 2 7 Neon 2 8 Sodium 2 8 1 Magnesium 2 8 2 Aluminium 2 8 3 Silicon 2 8 4 Phosphorus 2 8 5 Sulphur 2 8 6 Chlorine 2 8 7 Argon 2 8 8 Potassium 2 8 8 1 Calcium 2 8 8 2 7 Li 3 Nucleon number Atomic number
(See figure 25 for a diagram of some elements)
ISOTOPES
Are atoms of the same element which have different numbers of neutrons.
Most of the time isotopes have the same chemical properties (electrons usually determine an elements chemistry. Neutrons just give mass)
It changes the mass and the density though
Isotopes are classed according to whether they are stable or unstable (radioactive)
All chlorine atoms have 17 protons. Some chlorine atoms have 18 neutrons giving them a mass of 25. Some chlorine atoms have 20 neutrons giving them a mass of 37These atoms are called isotopes of chlorine.
Isotopes are atoms that have the same number of protons but different numbers of neutrons in their nuclei.
14-Carbon can be used in dating old objects
Uranium-235 and Uranium-238 are isotopes. They are not chemically different though, because protons the same therefore electrons the same.
RELATIVE ATOMIC MASS
75% of the chlorine in existence is Cl-35 25% of the chlorine in existence is Cl-37
The atomic mass shown in the periodic table is an average of the masses of both isotopes taking into account their relative proportions
(0.75x35 + 0.25x37) =35.5
SEPARATING TECHNIQUES
Separating mixtures of solids: Evaporation, sieving, crystallisation, magnets Immiscible liquids: Separating funnel
Insoluble solids from liquids: Decanting, filtration, centrifugation
Separating solution: Fractional distillation (>2 liquids of different boiling points), simple distillation (2 liquids of different boiling points), chromatography
Filtration-Separate an insoluble solid from a liquid. The liquid produced is called the filtrate. The solid left is the residue.
Evaporation: Separate soluble solid from liquid. Usually you heat the liquid in an evaporating basin.
Crystallisation: Similar to evaporation, although the solvent (liquid) evaporates slowly so that crystals form
Distillation: used to get the liquid from a solution by boiling it and condensing the vapour Fractional distillation: used to separate liquids with different boiling points (e.g. crude oil). Normally requires a fractionating column and a Liebig condenser.
Chromatography: Used to separate similar liquids, such as ink. Other solvents can be used apart from water (such as alcohol), but the beaker has to be covered to prevent solvent evaporating Melting points (m.p.) and boiling points (b.p.):Pure substances have definite melting and boiling points (temperatures)
Soluble: Does dissolve
Solution: formed when a substance (solute) dissolves in a liquid (solvent) Think: Solute is soluble.
Centrifuge: A machine with rapidly rotating container that separates very small solid particles from the liquid (i.e. Blood into red blood cells and plasma)
Saturated solution: A solution which has the maximum amount of solute dissolved in it Immiscible: Cannot undergo mixing or blending (liquid)
Miscible: Able to be mixed (liquid)
Separating funnel: A piece of glassware used to separate two immiscible liquids of different densities
Locating agent: RADIOACTIVE ISOTOPES
Medical use. Biological cells are sensitive to radioactive emissions. Radioactive isotopes are used in radiotherapy to treat cancer
Industry use: Use of Uranium-235 in nuclear power stations. The isotopes split into smaller parts. As a result huge amounts of energy are released.
TOPIC 2 NOBLE GASES
1st shell maximum 2, 2nd shell maximum 8, 3rd shell maximum 8 Each shell has the maximum number of electrons that it can hold Noble gases have shells that are completely full rather than partially full
During bonding, other atoms try to attain the “full electron shell” structure of the noble gases Group 0=Inert gases=Noble gases
All the elements in this group have fully occupied outer shells
They are the least chemically reactive elements. The stable arrangement means that the noble gases do not form chemical bonds with each other or other elements.
They are monatomic which means they exist as individual atoms
Noble gases reactivity make them very useful-prevents undesirable reactions from occurring (e.g. air ships can‟t be blown up easily if have helium. But if they were filled with hydrogen then they could very easily be blown up because hydrogen is very reactive but helium is a noble gas) Also, reactions can‟t take place in an inert environment which can be advantageous
TRENDS AND PHYSICAL PROPERTIES
Density increases down the group (NB: The noble gas group in order from period 1 downwards is Helium, Neon, Argon, Krypton, Xenon, Radon)
He density of 0.17g/dm cubed (dm cubed is just litre), Ne=0.84, Ar=1.66, Kr=3.46, Xe=5.45, Rn=8.9 Boiling point increases down the group
He= -269 degrees C, Ne= -246 degrees C, Ar= -186 degrees C, Kr= -152 degrees C, Xe= -108 degrees C, Rn= -62 degrees C
Noble
gas Uses
Helium Balloons big and small (e.g. air ships), protective gas for growing silicon crystals in silicon chip manufacture, rare documents (i.e. US Declaration of independence, prevents oxidisation), pressuring agent for liquid fuel rockets, coolant for the super conductors used in body scanners, inert atmosphere for welding (melting the metal to use as like a glue with blow torch)
Neon Fluorescent lights and „neon‟ signs, TV tubes (the big fat TVs), neon lasers
Argon Used to fill light bulbs because the filament will not react with argon, used for arc welding and cutting especially metals easily damaged by oxygen (e.g. aluminium & stainless steel) involved in dating the ages of rocks by potassium-argon dating
Krypton Used with argon for fluorescent lights, used in some photographic flash lamps, UV lasers
Xenon Used in making electron tubes, strobe lamps, high intensity lamps for projecting films Radon Used in earthquake prediction, used in radiation therapy in hospitals
EXOTHERMIC AND ENDOTHERMIC REACTIONS En in like enter. Ex is like exit
Energetics: the study of energy changes that occur in chemical reactions
When chemical reactions occur, along with new products forming there is usually heat change Products will have a different energy content than the original reactants
ENDOTHERMIC
If products contain more energy than the reactants, heat is absorbed or taken in by the surroundings and the change is called ENDOTHERMIC
The temperature of the system will fall in an endothermic change
Things that are cold take in heat energy (you are making bonds so energy is absorbed) Things that are hot give out heat energy
Energy changes in chemistry known as H
ENERGY LEVEL DIAGRAM FOR AN ENDOTHERMIC PROCESS
FIG b
Endothermic reactions are less common than exothermic reactions The products have more energy than the reactants (endothermic) Energy needed to overcome the attractive bonds
EXOTHERMIC
Products have less energy content than original reactants
If products contain less energy than the reactants, heat is released or given out to the surroundings and the change is called EXOTHERMIC
The temperature of the system will RISE in an exothermic change
Combustion reactions are always exothermic (the energy they release we use) Water is a substance low in chemical energy
Exothermic reactions convert chemical energy to heat energy
Products do not have more energy than the reactants in an explosion TYPES OF BONDING
Atoms can be joined together in three possible ways:
Metal and non metal (ONLY), Non-metals only, Metals only
Ionic Bonding Covalent Bonding Metallic Bonding
All three types involve changes in the electron structure in the outermost electron shells of the atoms
Most ionic compounds contain a metal and a non metal
When metals react they always lose outer shell electrons to leave a full electron shell This produces a charged atom (ion) with a positive (+) charge
FIG d
FIG e
FIG f
When non-metals react with a metal they gain electrons to achieve a full electron shell
The oppositely charged ions are attracted into a lattice (lattice structure=regular structure) that gets bigger until it consists of millions ions
Ionic structure-chemical bonds are caused by very strong electrostatic attractions between + and – ions
Metals lose electrons to form positive (+) ions called CATIONS Non-metals gain electrons to form negative (-) ions called ANIONS What is Sodium Chloride?
A compound made up of the two elements, sodium and chloride (they are chemically bonded (talking about electron movement))
Commonly known as salt An unreactive solid
Note: the elements that make up the compound have very different properties (sodium and chloride both unreactive, NaCl isn‟t)
All atoms want to have full shells in their structure because this makes them more stable and unreactive
E.g. Sodium (Na) has a „spare‟ electron in its outer shell and chlorine (Cl) has one missing electron from its outer shell, so sodium can transfer its „spare‟ electron to chlorine.
]
(Technically an atom can gain or lose electrons, but it likes the easiest route)
FIG g
Non-Metals
FIG h
FIG i
Fig j-Now both atoms have full shells. Electrons have a negative charge. Chlorine has gained an extra negative charge and sodium has lost a negative charge.
They are now known as IONS (charges particles)
Positive and negative ions are attracted to each other forming a strong ELECTROSTATIC BOND Na-Cl
Draw the dot and cross diagram for magnesium oxide. Don‟t forget the CHARGES on the newly formed ions.
The formula for magnesium oxide: MgO IONS:
POSTIVE IONS (CATIONS)
1+ 2+ 3+
NEGATIVE IONS (ANIONS)
1+ 2+ 3+
(SEE 2 PAGES OF REFILL FOR IONS DIAGRAMS) Fig k
FEATURES COMMON TO IONIC BONDING (SUMMARY OF IONIC BONDING) Metal atoms always use their outer electrons to form positive ions (cations)
Non-metals, with the exception of hydrogen, always gain electrons to become negative ions (anions)
The number of negative charges on a non-metal ion is equal to the number of electrons gained Ions formed are more stable
IONIC BONDS RESULT FROM ELECTROSTATIC ATTRACTION BETWEEN OPPOSITELY CHARGED IONS CREATING A LATTICE STRUCURE OF ALTERNATING POSITIVE AND NEGTIVE IONS
Because the lattice is rigid(Stiff and unmoving), this means that one gets a solid: ions don‟t move much
All atoms want to have a full outer shell COVALENT COMPOUNDS
They are formed when non-metal atoms react together e.g. Hydrogen
Their outer electrons are attracted to the positive nucleus and become shared by the atoms The shared electrons count towards the shells of both atoms and therefore fill up both shells (They don‟t create ions, these compounds are not charged)
They are held together by this SHARING OF ELECTRONS. A pair of shared electrons is called a COVALENT BOND
It can be drawn in 3 ways-A full bonding diagram, shown as a pair of electrons (xx) or just a line.
Covalent bonding in hydrogen and chlorine:
Both hydrogen (1) and chlorine (2.8.7) needs one more electron to get full outer shell Draw the electronic structure of hydrogen (cross) and chlorine (dots)
Using moly mods, build the following molecules: HCl, H2O, Cl2, H2, CH4, NH3 DIAGRAMS:
Fig l
MULTIPLE BONDS
Mostly electrons are shared as pairs
There are some compounds where they are shared in fours or even sixes This gives rise to single, double and triple covalent bonds
Again, each pair of electrons is often represented by a single line when doing simple diagrams of molecules
SINGLE BOND Fig o DOUBLE BOND TRIPLE BOND
Carbon has 4 valence electrons and needs 4 more electrons for full outer shell Hydrogen has 1 valence electrons and needs1 more electron for full outer shell Oxygen has 6 valence electrons and needs 2 more electrons for full outer shell
Nitrogen (2.8.5) has 5 valence electrons and needs 3 more electrons for full outer shell (SEE REFILL FOR TWO PAGES OF MULTIPLE BONDS DIAGRAMS)
MACROMOLECULES
SIMPLE MOLECULAR COMPOUNDS;
Many covalently bonded molecules exist as solids with low melting points.
These molecules are held together by WEAK INTERMOLECULAR FORCES and are easily broken down by heat
The molecules are then free to move but HAVE NO CHARGE (on the overall molecule) unlike ionic particles so does not conduct electricity.
ALLOTROPES
(SEE PAPER WITH PRINTED PICTURES OF DIAMOND AND GRAPHITE STRUCTURE)
Carbon atoms can also have many different arrangements (e.g. graphite and diamond) But they are both pure Carbon (C)
© There occur when millions of carbon atoms join together with a huge number of covalent bonds
They have similar chemical properties but very different physical properties such as hardness, conductivity, slipperiness (lubrication properties e.g. in engines), melting points and density. DIFFERENT ARRANGEMENTS OF THE SAME ELEMENT ARE CALLED ALLOTROPES
DIAMOND
Each carbon atom is covalently bonded to FOUR other atoms Tetrahedral structure
All valence electrons are involved in bonding Non-conductor of electricity
High melting point
Uses: Drill bits, cutting tools (because it‟s incredibly hard), wear resistant parts, jewellery (good reflective properties, lasts long)
GRAPHITE
Each carbon atom is covalently bonded to THREE other atoms One valence electron is not involved in bonding
This free electron forms the weak attraction and can carry charge Can conduct electricity
Uses: Lubricant, pencil lead (weak intermolecular force in the layers allowing them to slide over each other), electrodes (A conductor where electricity enters or leaves an object)
Note:
Diamond is such a hard material because every atom is covalently bonded to four other atoms.
Graphite feels slippery because carbon atoms in graphite are covalently bonded to 3 other atoms. Weak attractions between layers allow them to slide.
GIANT COVALENT STRUCTURES: Silicon dioxide
Sand is an impure form of Silicon Dioxide (Sio2)
Although it is a compound, it has a giant covalent structure with certain similarities to diamond. Notice there are repeating units of SiO2
The atoms are arranged in a 3 dimensional tetrahedral structure
Each of the covalent bonds is very strong and requires a lot of energy to break Very high melting point
METALLIC BONDING
Describe metallic bonding as a lattice of positive ions in a „sea of electrons‟ Describe electrical conductivity and malleability in metals
Lattice: regular arrangement
Metal atoms form a giant lattice structure similar to ionic compounds
The outermost electrons on each metal are free to move throughout the structure and form a „sea of electrons‟ or delocalised electrons.
Having released electrons in this „sea‟ the metal atoms are left with a + charge. Metallic bonding is the attraction of + metal ions to the „sea of electrons‟
PHYSICAL PROPERTIES
Metals are not brittle. They can be easily deferred or shaped. They are malleable. Metal atoms are the same and exist in simple structures.
If something hits the substance, it simply moves to the next layer along
Metals can have their shape changed relatively easily Malleable-can be hammered into sheets
Ductile-can be drawn out into wires
As the metal is beaten into another shape, the delocalised electron cloud continues to bind the “ions” together
Some metals such as gold can be hammered into sheets thin enough to be translucent Metals conduct electricity
In metals the delocalised electrons can move anywhere throughout the structure The attractive forces between the positive ions have no set direction
When a voltage is applied the electrons line up and „carry the charge‟
Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout its structure. Electrons attracted to the positive end are replaced by those entering from the negative end.
Electrons free to move Mobile electron cloud allows the conductivity of electricity Fig s
The strength of the metallic bond increases as:
The charge on the positive ions increase (Al^3+ stronger than Na^+) The size (ionic radius) of the cation decreases (+ve ion)
The number of mobile electrons per atom increases
E.g. Al has a strong metallic bond due to having small size, and a 3+ charge with 3 mobile (movable) electrons per atom.
Note:
Features of graphite structure that means it can conduct electricity: Free electrons between layers of carbon atoms
Giant covalent molecules have the physical property of high melting and boiling points
Note:
Ionic-Conducts when liquid but NOT solid (e.g. Sodium Chloride-Salt, when molten will have charged particles that can move in a liquid. It can‟t move in a solid lattice structure)
Metals- Conduct as both solid and liquid
Giant covalent-Will not conduct and hard to melt PROPERTIES AND STRUCTURE
To know structure and bonding relates to physical properties DENSITY
The density of substances depends on how closely the atoms are packed together Giant structures, metals especially, tend to be dense because all atoms/ions are pulled tightly together
Small molecules often have lower densities
SOLUBILITY
Generally substances with giant structures do not dissolve easily (although many ionic compounds dissolve for a special reason)
This is because in giant structures separating the particles involves breaking chemical bonds.
“Like Dissolves Like”
Generally substances with charged particles (ionic) or POLARCOVALENT MOLECULES (molecule with a slight positive and negative end) e.g. dissolve in solvents with „charged‟ species, called polar solvents (e.g. water, alcohol)
IONIC SUBSTANCES dissolve in water
Substances with uncharged particles (non-polar covalent (e.g. oil) dissolve in solvents with „uncharged‟ species, called non-polar solvents (e.g. petrol, tetrachloromethene/CCl4) Giant structures, such as metals, diamond, graphite and silica rarely dissolve.
Because of the δ (delta=very little) positive and negative charges of hydrogen (δ+) and oxygen (δ-), it breaks apart Ionic compounds and they dissolve.
WATER IS POLAR-it has got slight positive and negative ends
OIL IS NOT POLAR-won’t dissolve in water, but will dissolve in petrol. Note:
To tell if something is IONIC or not, it can conduct when molten, but not as a solid. Fig t
Fig u
The one test to determine if something is metal or not, is if it will conduct as a solid (only metals can)
The point of seeing if something is soluble in petrol, then you are looking at the covalent side of things.
All metals are giant structures CONDUCTIVITY
Covalent substances do not conduct electricity
This is because in covalent substances the outer electrons are fixed (localised) between specific ions (electricity will occur because and when electrons and ions can move) Metals conduct electricity
In metals the electrons can, given a potential, move anywhere throughout the structure.
Ionic substances do not conduct electricity when they are solid
When molten or dissolved they will conduct (and undergo electrolysis)
This is because the electricity is carried through the solution by the ions which are free to move when the ionic compounds is molten or in solution
BRITTLE VS MALLEABILITY
Ionic compounds are very brittle
Opposite charges attract, so neighbouring ions are pulled together
When something hits the substance a layer of ions will be pushed so that they are next to ions with the same charge.
Metals are NOT brittle
The metal atoms are the same and exist in simple structures
If something hits the substance, it simply moves to the next layer along. Metals are malleable
MELTING AND BOILING POINTS
Generally substances with giant structures have high melting points and boiling points Small molecules have melting points and boiling points that increases as the size of the molecule increases Fig w Fig w Fig x Fig x Fig z
BOILING POINTS AND VOLATILITY
Small molecules have low boiling points
They turn into a gas more readily, so they are said to be VOLATILE
[Volatility-how easily a liquid turns into a gas-less energy needed to break bonds (e.g. nail polish is very volatile, it quickly evaporates and dries)]
When a simple covalent structure boils, the weak forces between molecules break. Little energy is needed to break these forces, so the temperature needed is low.
ELECTRON ARRANGEMENT IN THE PERIODIC TABLE
The electron arrangements of atoms are linked to position in the periodic table
Elements in the same group have the same number of electrons in their shell (e.g. K is in group 1, so 1 e- in outer shell)
For the main group of elements, the number of the group is the number of electrons in the outer shell
The period also has numbers. This number shows us how many shells of electrons the atom has. E.g. period 2 elements have 2 shells e.g. Li (2.1)
General trends in the periodic table (leaving out noble gases)
Fig ab
Note: There are 7 periods in the periodic table, and 8 groups in the periodic table. The biggest group in the periodic table at room temperature is actually solids (not metals: P)
GROUP ONE-ALKALI METALS
To know the trend of physical properties for the Alkali metals REACTIVITY OF METALS
Generally, the further to the left and the further down the table, the more reactive the metal is (think negative vectors)
All of group one is more reactive than group 2 e.g. Sodium is more reactive than Calcium.
Lithium, sodium and potassium are all less dense than water and so will float. Densities follow a general, although not a perfect trend.
ELEMENT SYMBOL DENSITY
Lithium Li 0.53
Sodium Na 0.97
Potassium K 0.86
Rubidium Rb 1.53
Caesium Cs 1.88
Density increases, as the group descends.
The atoms in the Group 1 elements are bonded together using just one outer shell electron per atom.
As a result, melting points are low compared to most metals.
ELEMENT Melting Point
Lithium 181
Sodium 98
Potassium 63
Rubidium 39
Caesium 29
Melting point goes down, as you go down the group. Trends in chemical reactivity
Reactivity increases down the group
Reactions all involve the loss of the outer electron which changes the metal atom into a metal 1+ ion
Losing this electron seems to get easier as we go down the group Reactivity and Electron structure
1. The outer electron (-) gets further from the nucleus (+) as you go down the group. This reduces the force of attraction.
2. The inner shells “shield” the outermost electron from the attraction from the nucleus Both factors make it easier to lose the outer electron as you go down the group
(Note: The two factors are DIFFERENT) Fig ac
NOTE: THESE SPECIFIC OFFICIAL NUMERIC VALUES DO NOT HAVE DO BE REMEMBERED
Reaction with water
The group 1 elements all react vigorously with water (they react more vigorously as you go down the group)
Hydrogen gas is produced which sometimes catches fire
An alkali is left behind in the solution, which is why these elements are often called \The Alkali Metals‟
N.B. Stored under oil to prevent them reacting with water (and moisture in the air etc) Reaction of Lithium
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