Properties of matter Macroscopic - Microscopic – Not Observed with observed with senses sensesStrength Atom that bondsThermal and electricalconductivity consists ofMagnetic properties Kinds of bondsSolubility Forces betweenMalleability and ductility bondsDensityMP and BP
Properties of mixtures The substances keep their original properties. The substances don’t have to be mixed in a fixed ratio. The mixture can be separated with simple methods.
Two mixturesHomogenous Heterogenous There are more than one phase.The substances are all Eg sand and waterin one phase Suspension – solids that float in liquids.Eg air Eg muddy waterYou cannot identify the Emulsion – steady mix ofdifferent components insoluble substances in aof the mixture. liquid. Eg milk, mayonnaise
Separation methodsBecause mixtures retain the originalproperties, we can separate the differentsubstances by physical methods
Pure substancePure substances are made of only onesubstance or a compound.
Pure substancesElements Compounds A compound consist of two orAn element only consists more different elementsof the same atoms. bonded together.An element can’t be Can only be separated byseparated into simpler chemical methods.substances. Compounds’ properties differElements are categorised from the individual elements.as metals, non-metals and Joined in fixed ratios.semi-metals. Chemical reactions occur during formation.
Formulae H2O is the formula of water O H 2O H H2 hydrogen atoms1 oxygen atom CaO, CaSO4, Ca(OH)2, NH4NO3, CO2, NH3 Diatomic molecules Ionic bonds
Ionic bondsGroup 1, 2, 6, 7Transition metals – give off 1, 2, or 3 electronsMono-atomic anion – ‘ide’Polyatomic ion/radical
Physical differencesMetals Non-metalsMetallic lustre Dull (except graphite &Electrical conductors diamonds)Thermal conductors Poor electrical andOpaque thermal conductorsMalleable and ductile Some solids = opaque,Solids @ room gases are translucenttemperature, except Hg Solids = brittleHigh MP and BP Low MP and BP
Semi-metalsGenerally haveproperties of metals,but a few non-metalproperties as well.Ability to conductelectricity increaseswith heat. (in contrastwith metals)
Properties of semi-metalsShiny or dullConduct heat and electricity better than non-metals, but weaker that metalsWhen heated, they can conduct electricitybetter.
Electrical conductors, Semi-conductors and insulatorsMetals – conductorsNon-metals – insulatorsSemi-metals – poor conductors, called semi-conductors
Thermal conductors orinsulatorsThermal conduction is the flow of heat from ahigh temperature to a low temperature.Metals – thermal conductorsNon-metals – thermal insulators. All materialsthat trap air = poor conductors, because air is apoor conductor.
Magnetic and non-magneticmaterialsFerromagneticelements – stronglyattracted to magnetsFe, Ni, Co, Alnico,ceramic (insulatingmagnets), magnetite
Phases of matter And the kinetic molecular theory
SOLIDSSolids keep its shape and can only be dented, broken orbent. HardHigh density No compressibilityFixed volumeMade up of small particlesVibrate onlyVery small spaces between particlesStrong attractive forces – causes specific shapeNo diffusionHave crystalline structureHave a specific melting point
LIQUIDSNo fixed form Not hardHigh density No compressibilityFlows Fixed volumeParticles move in ordered fashionCollisions occurDiffusion occurSmaller spaces between particles than with gasesExerts pressure in all directionsWeak force between particlesSpecific freezing point and boiling point
GASESNo fixed form Not hardLow density Easily compressibleFlows No fixed volumeParticles move fastGreater collisionBig open spacesWeak/no forces between particlesInvoluntary motionDiffusion occursExerts pressure in all directions
Kinetic model of matter1. All matter consist of small particles2. Particles are in constant motion3. Spaces between the particles4. Constant collisions between particles and container5. Temperature is a measure of the kinetic energy of the particles6. Forces between particles7. Phase changes occur when energy changes occur
1. CondensationPrior to condensation:* particles slow down* not far apart* less violent collisionsPhase change follows:* Spaces decrease* Forces increase* more orderly arrangement
2. SolidificationPrior to solidification:* particles move very slowly* particles very close to each other* only vibratesPhase change follows:* very small spaces between particles* forces between particles become very strong* orderly arrangement
3. MeltingPrior to melting:* particles move fast* particles further apartPhase change follows:* spaces increase* forces decrease* less orderly arrangement
4. EvaporationPrior to evaporation* particles move very fast* particles very far apart* violent collisions due to high speedPhase change follows:* spaces between particles are big* forces negligible* disorderly arrangement
Evaporation vs. boilingBOILING EVAPORATIONOccurs @ B.P. Occurs @ temp belowOccurs throughout the B.P.liquid Occurs only on surfaceQuicker SlowTemp remains constant Causes cooling → heat during boiling absorbed from environment
The electrical nature of matterMichael FaradayElectrical current through salt solutionsAmount of Q = amount of atoms reacting
Daltons atomic theoryMichael FaradayAmount of Q = amount of atoms reacting●
Thomson’s atomic modelCharge and mass of electrons of all - substances are the same. -Thus, electrons in all substances must be the - same.Substances differ - because electrons are arranged differently.
Rutherford continuedObservations α – (+2) Most of the α-particles showed no diversion. + Atom + + Some were deflected Some were reflected + + + +Conclusions Electrons occupy the + + + greatest volume of the atom. + + + + Positive particles are grouped together in the nucleus – very heavy, but small.
Assumptions1. The positive charges are all together in a small volume in the nucleus.2. The nucleus is surrounded by a space that contains the e- (v. Small mass) – e- are responsible for the great volume of an atom.3. Mass is concentrated in the nucleus.Later investigations predicted that the nucleus ispositively charged. # Protons = # electrons. e-don’t move like bees around a hive, e- would collapseinto nucleus.
Bohr’s atomic modelElectrons move in orbitsElectrons with the same energy move aroundin the same orbitElectrons in orbits further away from nucleushave a higher energy
Planetary atomic modele- move in energy levelse- with same E values, move in same E levelsValence orbitals have higher energy than those close to the nucleusEnergy levels closer to the nucleus are filled first with e-Each energy level can only take a specific amount of e-e- in orbits close to the nucleus are lower in energy level than • When e- absorb energy orbits further away it rises in energy level.If e- are in lowest possible • This (excited) state is energy level – ground state unstable and e- fall back to lower energy levels
Line spectraElectrons in the ground state absorb energy. Electrons are now excited, and move to a higher energy level.This electron is now unstable.It falls back to its ground state, radiating extra energy as light.The separate coloured lines of light show electrons only have certain energy. Electron’s energy is thus quantised.Each element has it’s own unique line spectrum.Line spectra occur when gases are heated of an electric current is passed through it.
Wave mechanical atomic modelBohr’s atomic model explains the structure of hydrogen, but not those of atoms with more than one electron.The discovery of wave properties of electrons gave us a more acceptable model.e- have both particle and wave properties.Schrodinger stated that moving e- form a 3D wave space that surrounds the nucleus, called an orbital.
NeutronJ. Chadwick discovered a particle with a mass nearly equal to the proton.Neutral charges, called neutrons.
Atomic mass and diameterAtoms are extremely small with small masses. ELEMENT AVG ATOMIC MASS Hydrogen 1,673 55 X 10 -27 Carbon 1,994 36 X 10 -26 Oxygen 2,656 59 X 10 -26 Uranium 3,952 33 X 10 -25Diameters are also extremely small.Most of the volume of an atom is empty space, thenucleus accounts for most of the mass of an atom atthe centre.
Relative atomic massHydrogen is the lightest atom and is chosen as the standard for anatomic mass scale.This mass is equal to 1.Using proportion to find the atomic masses of other elementsrelative to a mass of 1.1,673 55 X 10 -27 kg of Hydrogen = 1 on the Hydrogen scaleSo, 2,656 59 X 10 -26 of oxygen = 1x2,656 59 X 10 -26 kg 1,673 55 X 10 -27 kg = 15, 87 on the Hydrogen scale
Structure of the atomThe atom consists of very many small subatomic particles.In chemistry we work with protons, neutrons and electrons.Protons and neutrons are in the nucleus at the centre.Electrons occupy a large region around the nucleus and are 1836times lighter.When electrons and protons are equal in number, the atom isneutral.When an electron is removed, the atom will be positivelycharged. An ion is an atom that has a charge on it.Cations are positively charged atoms.Anions are negatively atoms.
IsotopesSame element, different masses and amountof neutrons.Nuclide = isotopic nuclei.
Isotopes Carbon (atomic # 6) has three natural isotopes with atomic weights of 12, 13 and 14. isotope #p #n ====== == == C-12 6 6 C-13 6 7 C-14 6 8Tin (Sn, atomic # 50) has ten natural isotopes withatomic masses of 112, 114, 115, 116, 117, 118, 119,120, 122 and 124. How many protons and neutrons do these isotopes have?
Radioactive or Stable? Radioactivity is a nuclear phenomenon: it comes as a result of a particular structure in a nucleus. A radioactive atom is considered unstable. All unstable atoms emit radioactivity (usually by ejecting nuclear particles) in order to reach a stable configuration. This is the process of radioactive decaySo, not all atoms will be radioactive, just a small proportion of isotopes with unstable nuclei. The bulk of isotopes are stable, or non-radioactive.
Stable and Radioactive Isotopes Carbon (atomic # 6) has three natural isotopes with atomic weights of 12, 13 and 14. isotope #p #n ====== == == C-12 6 6 C-13 6 7 C-14 6 8C-14 is a radioactive isotope; C-12 and C-13 are stable. Over time the proportion of C-12/C-14 and C-13/C-14 will increase until there is no C-14. (unless some process makes new C-14...)
Radioactivity Inside You Concerned about radioactivity in nature?To keep things in perspective, consider that 0.01% of all potassium is radioactive K-40.Potassium is an essential element in the human body. If your body is about 1% K, this means a 70 kg (150 pound) person contains around 1x1021 atoms (that’s one billion trillion atoms) of radioactive K-40.
Energy levels in an atom Energy of an atom is quantized, meaning your electrons all have discrete amounts of energy.Electrons are thus limited to a specific energy level, Which you learned in grade 8 was an orbital. These main energy levels are indicated by n, and 1, 2, or 3 following it. eg. n = 1
Energy levels in an atom Each of these main energy levels are thensub-divided into sub-energy levels, which are indicated by the numbers s, p, d and f. s<p<d energy
Energy levels in an atomElectrons are constantly moving, and it is impossible to determine the position of an electron.Experimentation has indicated the most likely area of motion. This area is known as an orbital. s-level = one s-orbital, two electrons p-level = 3 p-orbitals, six electrons d-level = 5 d-orbitals, 10 electrons
Electron occupation of orbitalsThe distribution of electrons occur according to the Aufbau principle.●Paulis exclusion principle: An orbital can carry a max. of 2 electrons, ifit spins in opposite directions.●Electrons fill orbitals with the lowest energyfirst. The atom is most stable in its loweststate of energy.
Electron occupation of orbitals●When the same kind of orbitals areavailable, you first add one electron to eachorbital, and then fill them up. HUNDS RULE●2n² = amount of electrons in that orbital
ELECTRON CONFIGURATIONWith electronic configuration elements are representednumerically by the number of electrons in their shellsand number of shells. For example; Nitrogen configuration = 2 , 5 7 N 2 in 1st shell 2 + 5 = 7 5 in 2 nd shell 14
ELECTRON CONFIGURATIONWrite the electronic configuration for the followingelements; 20 11 8a) Ca b) Na c) O 23 16 40 2,8,8,2 2,8,1 2,6 17 14 5d) Cl e) Si f) B 11 35 28 2,8,7 2,8,4 2,3
sp-notation20 Ca: 1s²2s²2p63s²3p64s²20 Ca : [Ar] 4220 Ca 2+ : 1s22s22p63s23p6
Structure of the Periodic TableGROUPVertical columns.PERIODSSeven horisontal rows.VALENCE e- AND GROUP NUMBERValence e- are the same amount as thegroup number.
Structure of the Periodic TablePERIOD NUMBER AND ENERGY LEVELSThe period number indicates the energy levelin which the last e- are found.
Periodicity of the elementsPeriodicity is the recurring pattern of physicaland chemical properties as you move acrossthe Periodic Table.
Periodic LawThe elements of the same group have similarproperties. These properties differ from leftto right.
Physical and chemical properties of theelements are related to their atomic stucutre.
Atomic radius...is the distance from the nucleus and theoutermost stable electron orbital.
Atomic radius left to right in a period Atomic # increases, energy levels remainthe same. Electrons that are added are found in thesame energy level. Nuclear charge increases from left to right. Greater nuclear charges causes the e- to bereally attracted to the nucleus, causing theradius to become smaller.
Atomic radius top to bottom in a group # energy levels increase from top to bottom. Valence e- are further away from the core. e- in the inner energy levels shield the outere- (screening effect) Attractive forces of the core on the outer e-decrease. Atom volumes increase, so radius increase.
DensityForces of attraction and nature of substance determines density.Period – Period – non Groupmetals metalsincrease decrease Metals -decrease Non-metals - increaseMetal bonds Giant covalent Atomic massIncreased valence e-, structures.stronger forces, and Regular repeating increasestighter the atoms are pattern. Not packedpacked. closely together.
Ionisation energyFirst ionisation energy: energy needed toremove the 1st electron from a neutral atom inthe gaseous phase.X(g) + ionisation energy → X+(g) + e-
Factors influencing ionisation energy1. Charge of the nucleus2. Atomic radius3. Shielding effect4. Repulsion forces between e- pairs
Factors influencing ionisation energy1. Charge of the nucleusThe greater the atomic number, the greaterthe + charge, the greater the +/- attraction.
Ionisation energy - periodicityIonisation energy from top to bottom in a group. Usually decreases. Outer electrons are further away from nucleus, shielded by inner electrons.Increase in size, decrease in ionisation energy.
Consecutive Ionisation energyThe remaining electrons are attracted to the nucleus much stronger. e- repulsion decreases. Ionisation energy increases.
Electron affinityChange in energy when an electron is added to a neutral atom or ion in the gaseous phase.
Electron affinity Becomes more negative L → R. L → R: electrons are more stronglyattracted to the nucleus. Metals → low electron affinity. Dont take e-easily. Halogens → highest electron affinity. Takee- easily. Noble gases exempt – stable octets do notaccept e-.
Electron affinity Decreases from top to bottom. e- are further from the nucleus, weakerforce of attraction from the nucleus, electronaffinity decreases.
Electron negativity The amount of energy released when anelectron is added to an atoms in the groundstate, is called electron negativity.
Electron negativity Pauling scale – numbers on your periodictable. No units Increases from L → R in a period.Increased radius – e- are not tightly held. Decreases from the top → bottom in agroup. e- are held tighter by the nucleus.
Formulas of oxides When oxygen reacts with another element.Metal oxides are typical crystalline ionic solids. High M.P. And B.P. Dissolves in water – form bases.
Formulas of oxides Ratios: Metals –increases in each periodNon-metals – decrease in each periodNon-metal oxides – simple covalent that dissolves in water to form acids. Noble gases do not form oxides.
Formulas of halides React with metals – form ionic solids.React with non-metals – simple covalent molecules.Bonding ratio increases with metals and decreases with non-metals.
Similarities in chemical propertiesRead through p 121 – 124, revising the content from your gr 8 and 9 syllabus with your newly acquired knowledge.
Intramolecular forcesValence electrons are responsible for bonding between different atoms. I.F. explains the following:How thousands of atoms stay together in a molecule. The microscopic properties: M.P., hardness, conduction &c.
Intramolecular forces Covalent bonds, ionic bonds and metallic bonds
Covalent bondBond between non-metal and non- metal. Electrons are SHARED. Forms molecules
Relative formula mass Add the Ar of each atom.1. Formula of ionic compound2. Work out how many ions are present3. Ar(element 1) + Ar (element 2)4. No units
Metallic bonds1. empty/half-filled valence orbitals2. little energy is needed to loosen the valenceelectrons – delocalised3. atoms packed tightly – forms orderly crystallattice4. valence orbital overlap5. atoms positive charge – positive atomicresidue6. electrostatic force between atomic residueand sea of delocalised e-: metallic bond.
Metallic propertiesPROPERTIES REASONMetal glow Delocalised e- can reflect light – causes shiny surfaceConduction of Delocalised e- can move freely, actselectricity – solid/liquid as charge carriersConduction of heat Delocalised e- can move freely, act as heat carriersMalleability Atoms can slide over each otherHigh density Atoms packed closely together in a metal lattice
AlloysAlloys change a metals properties. They areusually always stronger and harder than thepure metal.STEEL – iron and carbonSTAINLESS STEEL – iron, chromium andnickelBRONZE – copper and tin