The Chemical Earth

1. The Chemical Earth

2. Why study chemistry? The Earth includes the biosphere, lithosphere, hydrosphere and atmosphere. Each of these is a mixture of thousands of different substances, many of which are useful to us if we: • have an understanding of the properties of the elements and compounds that make up the Earth’s materials • develop efficient processes for separating useful materials

4. Classification of matter • Pure substances have a fixed composition and fixed properties. They cannot be decomposed by simple physical separation techniques. • Mixtures have variable composition and variable properties. They can be separated into their components by various physical separation techniques.

5. Elements Pure substances can be further classified into elements and compounds. • Elements are the simplest pure substances consisting of only one type of atom. They cannot be broken down (or decomposed).

6. Compounds • Compounds are also pure substances. They are composed of two or more elements that are chemically bonded together. They are composed of a fixed number of atoms of each component element. They can be decomposed into their component elements or into simpler compounds.

7. Mixtures • The particles of each component in homogeneous mixtures are distributed uniformly (eg: sugar dissolved in water). • The particles in heterogeneous mixtures are not distributed uniformly (eg: concrete)

8. Chemical analysis Chemical analysis is the process of finding out what is present in a particular chemical sample. Chemical analysis can be: • Qualitative – determining what substances are present • Quantitative – determining how much of each substance is present in a sample.

9. Gravimetric analysis Gravimetric analysis involves separating the components of a material and accurately determining their mass. The percentage composition of the material can then be calculated. Gravimetric analysis can be used to determine the: • composition of a mixture using physical separation techniques • % composition of a compound using chemical and physical separation techniques.

10. Elements Elements consist of atoms of the same type. Many elements exist in nature as molecules. • Monatomic molecules: consist of only one atom (eg: noble gases) • Diatomic molecules: a molecule in which two atoms are bonded together. • Polyatomic molecules: a molecule of more than two atoms bonded together

11. Reactivity of elements Elements vary in their tendency to react. Elements that react readily (eg: calcium, sodium) are usually found combined with other elements as compounds. On the other hand elements such as gold which have low reactivity are often found in their pure form.

12. Metals • Relatively high densities • Good conductors of heat and electricity • Malleable and ductile • Shiny surface when freshly cut of cleaned • Relatively high melting points

13. Non-metals • State and form is variable (oxygen is a gas, bromine a liquid, sulfur a solid) • Usually not lustrous • Poor conductors heat and electricity • Not malleable or ductile • Variable melting points

14. Atomic theory An atom is made up of three fundamental particles: • Protons: positively charged particle • Neutrons: neutral • Electrons: negatively charged particle In an uncharged atom: number of protons = number of electrons

15. The nucleus • Contains protons and neutrons. • Has a positive charge equal to the number of protons • Contains about 99.9% of the mass of the atom • Extremely dense • Electrons move in space outside the nucleus

16. Atomic number and atomic mass • Atomic number (Z): number of protons in the nucleus and is fixed for any one element. (NB: In an uncharged atom this will also be the number of electrons) • Atomic mass (A): sum of the number of protons and neutrons in the nucleus

17. Atomic number and atomic mass Atomic mass = atomic number + number of neutrons OR A = Z + number of neutrons Therefore: A - Z = number of neutrons

18. A X Z Atomic number and atomic mass X is the element symbol A is the atomic mass Z is the atomic number

19. Electrons Electrons are believed to exist in energy levels or shells. The maximum number of electrons in each shell is determined by the formula 2n2 (where n = shell number 1, 2, 3, 4, etc). So: • 1st shell: can hold a maximum of 2 electrons • 2nd shell: can hold a maximum of 8 electrons • 3rd shell: can hold a maximum of 18 electrons • 4th shell: can hold a maximum of 32 electrons

20. Electron configuration The pattern of electrons in each shell is called the electron configuration. When determining the electron configuration of an atom the general rule is: Starting from the innermost shell, each electron shell or energy level must be filled before moving to the next energy level or shell. NB: potassium and calcium are exceptions

21. 23 Na 11 Sodium This means an atom of Sodium has 11 protons, 11electrons and and 12 neutrons. Therefore electrons are arranged as: • First shell = 2 electrons • Second Shell = 8 electrons • Third shell = 1 electron

22. Valence energy level The outermost shell of an atom is referred to as the valence energy level. Similarly, the electrons that occupy the outermost shell are called valence electrons. In the periodic table elements with the same number of valence electrons occur in the same column or group.

23. Octet rule In general, atoms are most stable when they have 8 electrons in their outer-most shell. This accounts for the lack of reactivity of the noble gases.

24. Ions Elements can achieve stable electron configurations by losing or gaining electrons. In doing so they form ions. • Positively charged ions (cations) are formed when one or more electrons are removed from an atom. • Negatively charged ions (anions) are formed when an atom gains one or more electrons

25. Electron dot diagrams Electron dot diagrams are a way of showing the arrangement of valence electrons in atoms. For example: 7 valence electrons 6 valence electrons 1 valence electron 2 valence electrons

26. Electron dot diagrams and ionic bonds Electron dot diagrams can also be used to show the formation of ions. Question: why will these two ions be attracted to each other and form an ionic bond?

27. Polyatomic ions Polyatomic ions are groups of atoms bonded to one another that have a net positive or negative charge. The carbonate ion is an example of a polyatomic ion. Polyatomic ions often have the suffix “ate” or “ite”. The superscript 2- indicates that there are two more electrons than the total number of protons possessed by the four atoms CO2- 3 The subscript indicates that there are 3 oxygen atoms

28. Covalent bonds Covalent bonds are formed when adjacent atoms share electrons. For example, a chlorine atom has the electron configuration (2, 8, 7). Two chlorine atoms can combine to form a chlorine molecule Cl2 by sharing a pair of electrons (each atom contributes one electron).

29. http://www.gcsescience.com/Chlorine-Formation.gif

30. Electron dot diagrams and covalent bonds Molecules are a group of two or more atoms held together by covalent bonds. Bonds in which two electrons are shared are called single covalent bonds and can be represented by a line drawn between the atoms H–H.

31. Electron dot diagrams and covalent bonds The number of covalent bonds formed by an atom depends on the number of valence electrons. In forming the molecule O2, each oxygen accepts a share of two electrons from the other atom. Hence four electrons are shared by the two oxygen atoms. This is called a double covalent bond: O=O

32. Ionic or covalent? If one member of a pair of atoms wants to gain electrons while the other wants to lose electrons then the pair will form an ionic bond. If both members want to gain electrons then they will form covalent bonds. Question: • Will Group I elements tend to form ionic bonds or covalent bonds? • What about Group 6 elements?

33. Valency The valency of an element is a measure of its “combining power” (the number of bonds it can form). • When an element forms ionic compounds the valency is the charge the atom carries. Ex: Na+ = +1 valency • When an element forms covalent compounds valency is the number of covalent bonds the atom forms. Ex: water is H-O-H, the valency of O = 2 and H =1

34. Formula ionic compounds • Ionic compounds are electrically neutral – therefore the number of negative charges must equal the number of positive charges. Ex: NaCl – the numbers of sodium and chloride ions is equal. • If the charges on the ions are not equal then there will be more ions with the smaller charge. Ex: the compound formed between Ca2+ and Cl- is CaCl2(there are 2 Cl ions for each Ca ion)

35. Naming ionic compounds The following naming rules apply to ionic compounds: • The cation (positive ion) is named first • The anion (negative ion) is named second • The suffix ‘ide’ is added to the non-metal in simple binary compounds (compounds made up of only two elements) Ex: NaCl = sodium chloride

36. Formula covalent molecular compounds In covalent compounds the formula represents the number of atoms of each element in one molecule of the compound. This is also called the molecular formula. Ex: H2O – two atoms of hydrogen and one atom of oxygen

37. Naming molecular compounds • The name of the element closer to the bottom or left-hand side of the periodic table is written first. • The the suffix ‘-ide’ is added to the end of the name of the second element. • The number of atoms of each element is indicated by the prefixes ‘mono-’, ‘di-’, ‘tri-’, ‘tetra-’, ‘penta-’ or hexa-’, which stand for 1, 2, 3, 4, 5 and 6 respectively. NB: prefix ‘mono-’ is not used for the first-named element.

38. Chemical equations In a chemical reaction the arrangement of atoms is changed to produce new substances but atoms are neither destroyed or created eg: mass is conserved. Chemical reactions are represented by chemical equations.

39. Writing chemical equations • Reactants are on the left and products are on the right eg: magnesium + oxygen magnesium oxide • Coefficients written IN FRONT of formulas show the number of particles of that substance eg: 2Mg = 2 atoms of magnesium • Physical state of reactants and products is shown by (g), (l), (s) – gas, liquid, solid

40. Writing chemical equations • The number of atoms of each element must be the same on the left and right hand sides of the equation eg: 2Mg + O2 2MgO 2 x Mg 2 x Mg 2 x O 2 x O • The sum of the electrical charges on the left must equal that on the right

41. Balancing equations Balancing a chemical equation is done by changing the coefficients in frontof the formulas. • Write the word equation • Write the formula for all elements and compounds present • Alter the coefficients to balance the number of each type of atom on both sides of the equation • Write in the physical states

42. Decomposition reactions Decomposition is a chemical reaction in which a compound is broken down into their constituent elements or simpler compounds. This is achieved by adding energy as: • Heat (thermal decomposition) • Light • Electricity (electrolysis)

43. Synthesis reactions Synthesis is the process of forming a compound from its component elements or other compounds in a laboratory. It leads to the formation of a more complex substance. For example, ammonia can be synthesised directly by combining nitrogen and hydrogen gases at high temperatures and pressures.

44. Bond energy A chemical change generally involves the absorption or release of greater quantities of energy than a physical change. Reason: A chemical change involves the breaking of chemical bonds

45. Covalent & ionic bonding • In covalent bonds electrons are shared between atoms. The energy required to separate atoms joined by a covalent bond is referred to as the bond energy. • Ionic bonds are formed by the electrostatic attraction between oppositely charge ions. The energy required to break ionic bonds is referred to as the lattice energy.

46. Bonding and Physical Properties All substances are made up of atoms, molecules or ions. It is the organisation of these particles that determines the physical properties of a substance. Solids can be classified into four groups on the basis of their physical properties: • Ionic compounds • Covalent molecular compounds • Covalent network compounds • Metallic substances