Physical Sciences Grade 10 Chemistry Notes

PHYSICAL SCIENCES : GRADE 10 CHEMISTRY – is the Science of matter and the change it undergoes. CLASSIFICATION OF MATTER Matter: is anything that occupy space ( has Volume) and has mass. Substance: is something that consists of one type of matter. It has consistent composition...

PHYSICAL SCIENCES : GRADE 10 CHEMISTRY – is the Science of matter and the change it undergoes. CLASSIFICATION OF MATTER Matter: is anything that occupy space ( has Volume) and has mass. Substance: is something that consists of one type of matter. It has consistent composition and distinct properties. Material: is the matter needed to make things, all objects around us are made of materials. PROPERTIES OF MATERIALS. Strength - is the ability to withstand applied force without breaking, e.g. steel and carbon fibre are strong. Conductivity – ability of materials to allow heat and electric current to flow through them. For example: metals are good conductors and non- metals are insulators. BRITTLENESS, MALLEABILITY AND DUCTILITY Brittle materials will break or crack when force is applied e.g. glass. Malleable materials can be hammered into sheets e.g. most metals. ductile materials can be stretched into wires e.g. copper. MAGNETIC AND NON-MAGNETIC: magnetic materials can be attracted by magnets, these are materials containing Iron, Nickel and Cobalt, as well as their alloys. DENSITY the amount of matter ( mass) per unit volume. Less dense objects like Aluminium can float in water but more dense like Lead and Gold sink. MELTING AND BOILING POINTS These are temperatures at which substances change from solid to liquid and from liquid to gas respectively. The boiling and melting point of a substance tells us its state at room temperature. CLASSIFICATION OF MATTER : ndfvifb Using the properties described above, substances can be grouped (classified) info three different types: Elements, Compounds and Mixtures. MATTER MIXTURES PURE SUBSTANCES COMPOUNDS ELEMENTS HOMOGENEOU HETEROGENEOUS S MIXTURES MIXTURES ( Can not be split into ( Separated by chemical simple substance (Separated by physical methods) methods)  EXAMPLES : Classify materials used in everyday activities according to their properties by ticking appropriate columns Strong Conducts Brittle Malleable Magnetic High melting Heat and And point electricity Ductile Plastic Iron nail Copper Glue Newspaper Ceramic Plate 2. What makes objects listed above suitable for their functions?  MIXTURES: A mixture is a combination of two or more substances that retain their properties. There are two types of mixtures: 1.Homogeneous mixtures: A mixture with the same composition throughout. Examples are solutions like salt dissolved in water, air, beer and alloys. 2. Heterogeneous mixtures a mixture whose composition is not uniform is called a heterogeneous. For example, any substances the do not dissolve in each other. Common example of heterogeneous mixtures are: Suspension – Solid particles in a liquid Emulsion – Oil in Vinegar e.g. Salad dressing Gel – Liquid in trapped in solid e.g. Fruit jelly Fog – Liquid in gas e.g. Spray Paint Smoke – Solid in gas Foam – Gas trapped in liquid e.g. Shaving cream Solid foam – Gas trapped in solid e.g. Foam rubber EXAMPLE: Classify the following mixtures as homogeneous or heterogeneous. Cup of Coffee, Pizza , Fizzy cool drink, Stainless steel, petrol , brass , air , scrambled egg.  SOLUTION HOMOGENEOUS HETEROGENEOUS Coffee Pizza Stainless Steel Scrambled egg Petrol,Brass, air Fizzy Cool drink PURE SUBSTANCES -Are substances that consist of one type of substance only which can either be elements or compounds. -Elements are pure substances that cannot be broken down into simpler substances by chemical methods. They are made up of only one type of atoms. - Compounds are pure substances made of the atoms of elements that are chemically combined in fixed ratios. PROPERTIES OF MIXTURES, COMPOUNDS AND ELEMENTS Element Compound Mixture Composition Consist of one kind of Composition is Composition may atoms constant vary and consist of two elements or compounds Properties are unique Properties differ Properties are from those of same as those of Properties constituent constituent elements substance Cannot be separated Can be separated Can be separated by chemical or by chemical by physical Separation physical methods methods methods Models  ACTIVITY 1. Classify the following as pure substances or mixtures: Water , Tea , Salt water , Copper , Brass , Air , Oxygen. SOLUTION: Pure Substances Mixtures Water, Copper, Oxygen Tea, Salt water, Brass, Air CRITERIA FOR PURITY There are three ways that can be used to determine if a substance is pure or not. 1. MELTING POINT Melting point is a temperature at which a solid turns into a liquid. This is constant ( the same) for pure substance. 2. BOILING POINT Boiling point is a temperature at which a liquid turns into a gas when heated. It is unique for pure substances e.g. pure water boils at 100ºc. 3. CHROMATOGRAPHY Is a method used to separate the pure substances in a mixture. Different substances in a mixture separate out by dissolving in a solvent and bond with the absorbent material at different rates. They form different bands on the material. e.g.: Black ink is a mixture of different coloured compounds – If a black ink spot is placed on a paper strip ( chromatography paper) and the end dipped in solvent. The different coloured pigments would be observed after some time. During After Before Chromatography Solvent soaks up the papaer paper Black Ink component Colours of black Ink  Examples ACTIVITY 1. Caffeine is a compound found in coffee beans. The diagram below shows a chromatography of caffeine and four drinks, A to D.  caffeine A B C D 1. Which one of the drinks A to D contains caffeine only? Solution C 2. Chose an item from column B that matches a description / item in column A. Column A Column B 2.1 Potassium chloride crystals A heterogeneous mixture 2.2 A good conductor of electricity B homogeneous mixture 2.3 A non-metal element C silicon 2.4 Sugar dissolved in water D sulfur 2.5 Blood E compound 2.6 Increasing conductivity with increase in F copper temperature  ANSWER. 2.1 E 2.2 F 2.3 D 2.4 B 2.5 A 2.6 C ELEMENTS - Elements are made up of atoms of the same kind. - Elements can be represented in symbols of one or two letters, first letter is capitalised but the second letters are not, e.g. Co represent cobalt but CO is a formula for carbon monoxide, carbon and oxygen. COMPOUNDS - Compounds are pure substance that are made of two or more kinds of elements combined in fixed proportions. For example water is a compound of hydrogen and oxygen in the ratio of 2:1. - Each compound has its own formula that tells us: 1. The type of elements in a compound 2. The number of atoms of each element in the compound  Examples Name of compound Formula Atoms in the compound Water H 2O 2 hydrogen atoms, 1 oxygen atom Carbon dioxide CO2 1 carbon atom,2 oxygen atom Ammonia NH3 1 Nitrogen Atom, 3 hydrogen atom NAMING AND FORMULAE OF SUBSTANCE - When compounds are formed, metals form positive ions called cations and non-metals usually form negative ions called anions ( except for hydrogen, H+ and ammonium ions , NH4+). - First part of the name comes from the cation and second part from the anion. - The name ends with – ide for a single anion , - ite and – ate are used if oxygen is also present for example. 1. Fluoride F- -- as in sodium fluoride, NaF Chloride Cl- -- as in potassium chloride, KCl Sulfide S2- -- as in Iron Sulfide, FeS 2. Sulfate as in copper sulfate, CuSO4 Prefixes can be used to indicate number of same atoms in the compound: For example : Carbon monoxide CO - one Oxygen atom Carbon dioxide CO2 - two Oxygen atoms Sulfur trioxide SO3 - Three Oxygen atoms WRITTING CHEMICAL FORMULAE -Compounds have fixed chemical formula. The following can rules can be used in writhing a formula: 1. Write down the symbol of bonding ions, starting with the positive 2. Determine the charges of each 3. Interchange the charges , ignore signs, and write them as subscripts in front of symbols. -Example; 1. Sodium oxide Step 1 Na O Step 2 +1 -2 Na O step 3 Na2O is the formula for sodium dioxide  Aluminiun sulfate Step 1 Al SO4 Step 2 +3 -2 Al SO4 Step 3 Al2 (SO4 )3 is the formula. NB Charges are given in special tables for different ions at this stage.  METALS, METALLOIDS AND NON-METALS - Elements can be classified as metals, metalloids and non – metals, - METALS They have the following properties in common: - Found on the left hand side of the periodic table, groups 1,2, up to 13. - All are solids except mercury which is a liquid at room temperature - They are shiny, malleable, ductile and greyish except for gold and copper, - They are good conductors of heat and electricity.  Non – Metals They are found on the right–hand side of the periodic table, groups 14 to 18. They are either solids or gases at room temperature except bromine which is a liquid. They have low density and vary in colour. Non-metals are poor conductors of both heat and electricity except graphite. Most of them are brittle. METALLOIDS Metalloids are found between metals and non metals in the periodic table. They have mainly non metallic properties but show metallic properties under certain conditions. There are seven metalloids, Boron(B) , Silicon(Si ), Germanium(Ge), Antimony(Sb), Tellurium(Te), Arsenic(As), Polonium(Po).  ACTIVITY 1. Write down the uses of metals, metalloids and non-metal based on their properties. Solution Metals – are shinny and can be used as ornaments e.g. gold and silver Are strong and can be used in construction e.g. iron. Are good conductors of electricity and can be used as electrical cables e.g. copper NON – METALS – are poor conductors, so can be used as insulators. SEMI –METALS – Increase conductivity with rise in temperature , are used in electronic industry. MAGNETIC MATERIALS Magnetic materials are those which can be attracted to a magnet. There are three elements(metals) with magnetic properties; Iron, Cobalt and Nickel. Their alloys are also magnetic. The rest of the materials are non-magnetic Activity 1.List some uses of magnets and magnetic material in daily life Solution They are used in speakers, telephone, electric motors and generators , as compasses etc.  TOPIC 3:THE ATOM An atom is the basic building block of all matter. Identical atoms come together to form elements. THE ATOMIC MODEL A model is a representation of an object or a system. A model can be real or mental picture of a difficult concept or phenomena in terms of what we know. Models are used to explain certain observations. They are not stagnant but are modified and developed over time. Many Scientists came up with models of an atom that were developed into laws and theories over a period of time. A few to note are: 1. JOHN DALTON’S ATOMIC THEORY. ( 1799 – 1844) Elements are made up of tiny particles called atoms Atoms of a given element are identical Atoms of one elements is different from those of any other elements Atoms of one element can combine with atoms of another element to form Chemical compounds. Atoms cannot be created , divided into smaller particles or destroyed in the Chemical process.  2. J.J THONSON ( 1897 ) He pointed out that atoms have negatively charged electrons ( plums ) embedded in a positively charged ‘ pudding ‘. This model was called a plum pudding model. 3. ERNEST RUTHERFORD ( 1871 – 1937 ) -Using a gold foil and alpha particles ( Helium nuclei ) concluded that atoms have positive charges ( protons ) in the center ( nucleus ). The nucleus is surrounded by negative charges called electrons. -This model of an atom is still being used today 4. NEILS BOHR ( 1885 – 1962 ) -Suggested in his model that electrons revolve around the nucleus in orbits or shells called energy levels. -Orbits are several fixed distances from nucleus. - Electrons in each level have a specific amount of energy. 5. JAMES CHADWICK ( 1932 ) -He discovered the neutron , in the nucleus and with no charge. Examples - Some models are called laws and yet others are called theories. What is the difference between a theory and a law?  Answer A law is a theory that has stood up to a great a deal of testing(experimentally) and proven beyond doubt. Theory is a statement that just explains a fact but has not been experimentally proved. ATOMIC MASS AND DIAMETER Atoms are very, very small , so their mass and diameter are given special units. The smallest atom , hydrogen atom has a mass of 1,0078 atomic mass unit ( U ) and radius of 1,1 angstrom ( Å ) Note 1 U = 1,66 x 10 -27 Kg 1 Å = 1x10-10 m  Atoms consist of three particles; protons( P+) , neutrons(n0) and electrons ( e- ). The properties of the sub atomic particles can be summarised in a table as follows: 12  Number of neutrons = mas number–proton number n=A–Z For a neutral atom , number of electrons is equal to number of protons ( atomic number ). Examples Consider the boron atom. List number of protons, neutrons and electrons in a neutral boron atom and in a boron ion ( B3+ ) Answers Boron atom has 5 proton , 5 electrons and 6 neutrons Boron ion ( B3+ ) has lost 3 electrons , it has 5 protons 2 electrons and 6 neutrons.  ISOTOPES Isotopes are atoms of the same element with different number of neutrons(or mass number ). Isotopes have same atomic number , same number of electrons but different masses. chemical properties of isotopes are the same. The relative atomic mass of an element is sometimes not a whole number because it represents the average of two or more isotopes of the element.  Examples 1.Chlorine has two isotopes , chlorine –35 and chlorine-37. In a sample of chlorine 75.5% is chlorine-35 an 24.5% is chlorine -37 atoms. Calculate the relative atomic mass of chlorine.  Answer Note Chlorine–35 means chlorine of mass 35 R.A.M. = % abundance x mass of isotopes + % abundance x mass of isotope 100 = ( 75,5 X 35 ) + ( 24,5 X 37 ) 100 = 35,5  Example 2 Naturally nitrogen occurs as the two isotopes 14N and 15N, How many different types of nitrogen molecules will occur in air as a result of the combination of the above atoms?  Answer 1 molecule, the isotopes have same chemical properties so only one molecule is formed.  ELECTRON CONFIGURATION The atom’s electron configuration is the way in which electrons are arranged around the nucleus. Atomic orbitals Electrons move in orbits or shells that are called energy levels. The lowest energy level level(ground State) is found around the nucleus , electrons in this level are most stable. Energy levels are numbered 1, 2, 3, and so on. In each energy level there can be the orbitals called the s, p, d and f. s orbitals are spherical and can accommodate a maximum of 2 electrons. p orbitals are dumb – bell shaped , made up of there identical orientations in there dimensions; x , y and z direction so p has a maximum of 6 electrons (2 each); Px , Py and Pz d and f orbitals are for elements with energy levels 3 and above, this is not going to be covered at this level.  THE ELECTRON CONFIGURATION USING AUFBAU PRINCIPAL This configuration tells us how electrons are distributed in the various orbitals, it can be done in two ways; 1. The spectroscopic configuration notation( S.C.N.) Number of electrons in each orbital is shown as superscripts of the orbital e.g. sodium(Na) has 11 electrons , its configuration is: 1S2 , 2S2 , 2p6 , 3S1 2. The orbital box diagram Orbitals are represented by a circles or a boxes , electrons are represented by arrows. General rules in filling in the orbital box diagram 1. Use a block or a circle to represent an orbital. 2. Arrows and represent electrons in opposite Spin. 3. Lower energy levels are occupied first. 4. An orbital can have a maximum of two electrons in opposite spin, this is called Pauli’s Exclusion principle. 5. Single electrons occupy each orbital in the same energy level before they start to pair, this is called Hund’s rule. 6. Order of orbitals for the first 20 elements is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p 7. The maximum number that can be in a certain energy level is given by formula 2n2 , where n is the energy level number. 8. e.g. level 1 has 2x12 = 2 9. level 2 has 2x22 = 8 10. level 3 has 2x32 = 18 Note: these are maximum values, an energy level can have. EXAMPLES 1. Nitrogen (N) has 7 electrons because it is a neutral atoms configuration of N is 1s2, 2s2, 2p3 Aufgau’s orbital box diagram 2p 2s 1s can not pair before filling orbits in the same level 2. Oxygen (O) has 8 electrons because it is a neutral atom, its configuration is 1s2, 2s2, 2p4 2p 2s 1s 3. Sodium ion Na+ has 10 electrons , the atoms has lost 1 electron to become an ion Configuration: 1s2, 2s2, 2p6 Orbital box diagram : 2p 2s 1s 4. Chlorine ion, Cl- has 18 electrons, the atom gained an electron to become an anion. Configuration : 1s2, 2s2, 2p6, 3s2 3p6 Orbital box diagram 3p 3s 2p 2s 1s  CORE AND VALENCE ELECTRONS Valence electrons are electrons in the outermost energy level of an atom. Core electrons are electrons in the filled energy levels. If an element has only core electrons, then it is stable, has less potential to react than these with valence energy levels. This is the case with noble gases, group 18 elements.  TOPIC 4: THE PERIODIC TABLE The periodic table arrange elements in order of increasing atomic number and group them according to their physical and chemical properties. It is described in Groups and Periods: Group–are vertical columns with similar properties, they are number from 1 to 18. Periods - are horizontal rows, with same number of energy levels, periods, 1 to 7.  Some groups have specific names - Group 1 - alkali metals - Group 2 – alkali earth metals - Group 3 to 12 are called transition elements - Group 17 are halogens - Group 18 are noble gases Elements in the same group have same valence electrons, for example Group 2 :  Be : [He ]; 2s2 Mg : [Ne ]; 3s2 Ca : [Ar ]; 4s2 Each member has a full set of core electrons as a noble gas in previous periods, shown in [ ] and 2 electrons in the S orbital. The number of valence electrons determines the chemical properties because they are involved in chemical bonding. The number of electrons an atom transfer ( loose, gain or share) to achieve a stable octate (like noble gas) is called its valency. Valence and valency for each group are show below. Group 1 2 13 14 15 16 17 18 Valence electrons 1 2 3 4 5 6 7 8 Valency 1 2 3 4 3 2 1 0 Note: Number of valence electrons can also be used as group numbers if transition electrons are ignored, they are usually written in Roman numerals as I, II, III, IV, V, VI, VII VIII. PATTERNS IN THE PERIODIC TABLE 1.Density -Generally density increases to the middle of the period and decrease again to the right. This is because of the type of bonds between atoms. 2.Melting and Boiling Points -Like density they increase to the middle of the period then decrease to the right. Group 14 have highest melting points because they make macromolecules due to high valency. 3. Atomic radius - Increase as we go down each group , since number of electron energy levels also increase as we go down. - Decrease across the period from left to right because the increase in nuclear positive charge attract electrons closer and become smaller. IONISATION ENERGY Ionisation energy is the amount of energy required to remove an electron from an atom. First ionisation energy is energy required to removed the first electron from a neutral atom i.e. M + 1s+ ionisation energy M + + e- M+ + 2nd ionisation energy M2+ + e- - It is more difficult to remove an additional electron - First ionisation energy generally increase across the period. It is easy to ionise metals than non metals. ELECTRON AFFINITY - Electron affinity is the energy change that occurs when an atom accepts an electron in its gaseous state to form an anion. X ( g ) + e- X- + electron affinity. - It increase from left to right across the period - It is easy for non metals to accept electrons than metals. ELECTRONEGATIVITY - An ability of an atom in a compound to attract the bonding electrons to itself. - Electronegativity increase from left to right, group 17 elements are more electronegative than other groups  CHEMICAL PROPERTIES IN GROUPS Elements in the same group have the same valence electrons so they have the same chemical properties. GROUP 1 : The Alkali metals All have low density and float in water All are soft and can be cut by a knife Are silver, shiny on surface with fresh cut Are conductors of heat and electricity All burn in oxygen to form soluble oxides (in water) form alkaline solutions (hydroxides) that’s why they are called alkali metals. They form cations with charge +1 and reactivity increase down the group.  GROUP 2 :The alkali earth metals Silvery metals that form cations with a +2 charge and are good conductors of heat and electricity. Burn in air to form solid oxides which are partially soluble in water. GROUP 17 : The halogens Are non metals which form diatomic molecules Form anions which a –1 charge Have low melting point and low boiling point which increase down the group. All are coloured and the colour become deep down the group.. Their reactivity decrease down the group Group 18 : Noble gases All are extremely unreactive but reactivity increase down the group. For example Krypton and Xenon react slightly. Boiling points increase down the group. Examples 1. When an atom X of an element in, Group 1 reacts to become X+, which property of X has changed? Answer:X has lost an electron, so the number of filled energy levels decreased. 2. M is a yellowish gas at room temperature, it forms a charge -1. In which Group does it belong? Answer Group 17 , Halogens.  CHEMICAL BONDING Atoms react in order to result in a more stable product , that is similar to a configuration of noble gases. There are three types of reactions: 1. Metal and metal resulting in metallic bonding 2. Non – metal and non – metal resulting in covalent bonding 3. Metal and non – metal resulting in ionic bonding. Lewis dot diagrams Valence electrons are the only electrons involved in chemical bonding , so Lewis dot diagrams only use then to show bonding. Examples of Lewis dot symbols are:  H hydrogen Li Lithium Be Berylium B Boron F Fluorine and so on Note: Elements in same Group have similar Lewis dot symbols because they have same number of valence electrons Covalent Bonding form when two atoms share their electrons , a shared electron pair is represented by a line called a bond. For example : two hydrogen atoms H +H H H same as H H - Covalent bonding result in molecules. - We can use and X (dots and crosses) to distinguish between electrons of different atoms. For example: Hx + Cl Hx Cl An electron diagram can also be used as: + + X – Hydrogen --Chlorine H CL More than two atoms can also share electrons: For example in Ammonia molecule (NH3)  two electron pairs result in a double bond, and three in a triple bond. For example , 1. oxygen molecule: 2.) Nitrogen molecule has five valence electrons and will need three to be stable IONIC BONDING Ionic bonding involves transfer of one or more electrons from one atom to another forming ions. These ions are attracted to each other and pack into crystal lattices. -- Metals loose electrons to become positive charges called cations while non –metal gain electrons to become negative charges called anions.  For example In some cases the cation and anion may have different charges: Metallic Bonding -Atoms in a metal have low ionisation energy , so each atom loose control of its valence electrons and they become delocalised. -- valence electrons form a sea of electrons around positive atomic kernels. -For example:  Positive charges are ‘ glued ‘ by a sea of electrons. This model helps to explain why metals a good conductors, are malleable and ductile. Examples Information of two elements P and Q are given in the table below: Atomic number Mass number Electron Structure P 16 32 1s2 2s2 2p6 3s2 3p4 Q 3 7 1s2 2s1 a) Which element has two core electrons in each atom? b) Write down the formula of the compound formed by the combination of elements: P and Q c) What type of bonding is likely to take place between P and Q? Answers a) Q - element in the 1st energy level b) Q2P c) Ionic , P looses an electron and Q gains 2electrons.  Using electronegativity to decide whether bonding is covalent or ionic We find the difference in electronegativity between the two atoms in a bond.