Chemistry

2. Chemistry • Is the study of matter, its properties and its changes or transformations

3. Matter • Anything that has mass and takes up space

4. Physical Properties • a change in the size or form of a substance which does not change the chemical properties of the substance is called a physical change • a physical property is a characteristic of a substance • these properties allow us to distinguish or tell the difference between substances • examples of physical properties - state, color, odor, luster, texture, hardness, crystal form, mass, volume, density, solubility, viscosity, malleability, ductility, melting point, boiling point

5. Chemical Properties • is a characteristic behavior that occurs when a substance changes into a new substance\ • the change itself is called a chemical change • the starting materials we call reactants- the final materials we call products • examples of chemical properties are: reactions with water, reactions with acids, combustion • example:  iron + oxygen    =   rust

6. Introduction to Naming • Today, most compounds are known by their IUPAC names.  IUPAC stands for International Union of Pure and Applied Chemistry. This organization has determined a set of rules to be used for naming chemicals. Its purpose is to set international guidelines so that all scientists follow the same rules.

7. Classifying Compounds • There are three main types of compounds that we will be dealing with: • Ionic – a combination of metals and non-metals • Molecular or covalent – a combinations of two nonmetals • Inter-metallic or metallic – a combination of two metals

10. (cont’d) • Moleculesare combinations of two or more atoms. • Molecular element - if the atoms are all the same • For example oxygen gas is a molecule composed of two atoms of oxygen. Since there are two atoms the molecule is called a diatomic molecule. (just remember the gen's)

11. Diatomic molecular elements As the heading suggest these are elements composed of two (di) nonmetal  atoms.  We seen these in the last lesson. (just remember the gen's)

12. Mono-atomic elements If we look at the name of the heading, mono - means one, so these are the non-metals that that exist in nature as individual atoms.  Although these are not compounds we have included them here because we will reference them many times. Noble Gases (group 18)

13. Polyatomic molecular elements These are non-metal elements composed of many (poly) atoms.

14. Compound -a molecule that contains two or more different types of atoms or ions. • Water (H2O) is a compound because it contains both Hydrogen and Oxygen, two different types of atoms. • The formula for water (H2O) is a combination of symbols and subscripts. • H and O are the symbols for the two types of elements (Hydrogen and Oxygen) found in water. • The number 2 to the lower right hand corner of the symbol for Hydrogen is called a subscript.

15. a molecular formula indicates the total number of atoms  in one molecule • an empirical formula is the simplest whole number ratio of atoms in the compound. •  consider hydrogen peroxide (H2O2 ) as an example. • - the molecular formula is H2O2. • - the empirical formula is HO. (lowest ratio is 1:1) • In some cases the molecular formula and empirical formula are the same. • For example, both the molecular formula and the empirical formula for water is H2O. (the lowest whole number ratio is 2:1) • It is important to recognize however that the empirical formula only describes the ratio of one atom to another, and not the actual number of atoms of each type in the compound.

16. Solutions • Aqueous solutions are those in which the solvent is water. They form in at least three ways depending on the nature of the solute: • molecular solvation • dissociation • ionization

17. Electrolytes and Non-electrolytes • A solution that conducts electrical current is said to be electrolytic and the solute is called an electrolyte. • The solute in a solution that does not conduct electrical current is a non-electrolyte. • Generally, dissociated ionic compounds are electrolytes whereas dissolved molecular compounds are non-electrolytes. The exceptions to this rule are the molecular acids.  • Acids are defined theoretically as species that ionize in water to produce hydrogen ions and negative ions (anions).

18. Bohr’s Electron Energy Level Theory • Bohr was the first to attempt to describe the way electrons are distributed in an atom • The type of bonding that occurs in a substance is a function of the way the electrons are distributed in an atom

19. Electron Energy Level Theory • Electrons orbit the nucleus in specific pathways called energy levels • There is a fixed number of energy levels • Each energy level is capable of holding a certain number of electrons • The location of an electron (with respect to the nucleus) is an indication of the energy it contains (closer more – further away less)

20. Electrons can move from one energy level to the next by gaining or losing a specific amount of energy called a quantum • Energy levels are not equidistant from each other

21. The part of the energy level or region of space in which we find electrons are also called orbitals. These are not the fixed pathways we normally think of when discussing orbits, but rather they are a region or area when the electron would be located at some point in time.

22. Electron Energy Level Diagrams • These will demonstrate how electrons are distributed in an atom • Determine the atomic number of the element • Draw a circle to represent the nucleus and write the # of protons in the circle • Determine the period number of the element (ie what row it is in). This is equal to the number energy levels in the element

23. Each energy level can hold a specific number of electrons. We always fill the inner level first. • For the first three levels it is 2 -8 – 8. We put two hash marks - - for each level and put the number of electrons in the level between the marks. • The outermost or last energy level is called the valence level. 7. The electrons in the valence level are called valence electrons.

24. Example • Look up the atomic number of nitrogen and use it to determine nitrogen's number of protons and electrons. Since nitrogen is atomic number 7, it has 7 protons and 7 electrons. • Draw a circle to represent the nucleus and write in the number of protons using the shorthand 7p+. • Locate the period number for nitrogen and use it to determine the number of electron energy levels. The period number (2) equals the number of energy levels. Draw two broken lines above the circle to represent nitrogen's two electron energy levels.

26. Important Point!!! • Valence electrons are the only ones thought to be involved in chemical reactions!

27. Your Turn • Draw energy level diagrams for the following elements: • Lithium • Calcium • Neon • Hydrogen

28. -- 2 – -- 8 – -- 8 – -- 2 -- • - 1 -- • -- 2 -- -- 8 – -- 2 -- -- 1 -- 3+ 20+ 10+ 1+ Bohr Diagrams

29. Order of Filling Energy Levels • Lower energy levels are filled before higher energy level orbitals • The electrons will fill the available orbitals before pairing up.

30. Your Turn Construct modified Bohr diagrams for the first 20 elements

31. Why do certain atoms gain or lose electrons? • The loss or gain of electrons enables atoms to achieve an octet of electrons (ie. gain electrons) or expose a lower energy octet of electrons (ie. lose electrons)

32. Nobel Gases and the Octet Rule • Remember that the outer shell or energy level is called the valence level and it contains the valence electrons • It is commonly believed that the valence energy level and the valence electrons are responsible for chemical bonding • A group or eight is called an octet • The octet rule states that chemical stability is associated with a group of eight valence electrons • The noble gasses contain a complete outer shell of electrons or they have 8 valence electrons • The noble gasses are believed to be chemically stable because their valence level have a complement of eight electrons

34. Simple Ions • Simple ions are atoms that no longer have equal numbers of electrons and protons. • As a result of this imbalance they now have a positive or negative charge • Cation • A cation is  a positve (+) ion •  Metal atoms lose electrons to become positive thus, metals form cations.  • Example:  Na+, Mg 2+ • Anion • An anion is ANegative ION (A N ion) • Non-metal atoms gain electrons to become more negative thus, non-metals form anions:  • Example:   F- , O 2-

36. Your Turn Construct modified Bohr diagrams for the ions formed by the first 20 elements and name each ion formed Complete the simple ions sheet

37. Your Turn • Complete sheet u3 s2 l3 Simple Ions

38. Chemical Bonding • A chemical bond will form between two atoms if the attractive forces between two atoms is greater than the repulsive forces • Attractive forces are electrical forces that hold the atoms, ions, or molecules together • Bonds are formed through the valence electrons in the atom • Valence electrons are usually transferred from the outer shell of one atom to the outer shell of another atom or shared among the outer shell of combining atoms

39. Ionic vs Molecular Compounds • Ionic Compounds • Involve the transfer of electrons resulting in ionic bonding • Made up of two oppositely charged ions (metal and nonmetal) • Exist in the form or an ionic crystal lattice • Binary Molecular Compounds • Involve the sharing of electrons resulting in covalent bonding • Composed of two nonmetals • Exist as individual molecules

40. Molecular Compounds • Always form between two or more non-metals • In order to get full outer energy levels the atoms that make up molecular compounds share their valence electrons (bonding electrons) • If they dissolve in water, their solutions do not conduct electricity • Molecular compounds are non - electrolytes

41. Molecular Compounds (cont’d) • Individual parts of a molecular compound are called molecules. • The molecular formula tell us how many atoms of which elements are being shared. • Example: H2O tells us that in every molecule of water there are two hydrogen atoms and one oxygen atom

42. Naming Molecular Compounds (cont’d) • When we want to come up with (derive/make) a formula for a molecular compound there are two routes to follow as well: • 1. It has a trivial name and its formula must have been memorized, like water • 2. We use the prefixes given to tell us how many of each element, like silicon dioxide

43. Naming Molecular Compounds • Two basic Ways 1. the compound has a trivial (common) name that has been used for so long, it’s too late to change it. We must memorize trivial names and their formulas.

44. Trivial names, formulas, and Special Situations • The following names and formulas must be memorized • Water H20 • Ammonia NH3 • Glucose C6H12O6 • Sucrose C12H22O11 • Methane CH4 • Propane C3H8 • Octane C8H18 • Methanol CH3OH • Ethanol C2H5OH • Hydrogen Peroxide H2O2 • O3 ozone

45. Naming Molecular Compounds (cont’d) 2. We use a system of prefixes to signify how many of each type of atom is being shared to make up the molecules 1 = mono 6 = hexa 2 = di 7 = hepta 3 = tri 8 = octa 4 = tetra 9 = nona 5 = penta 10 = deca

46. Naming Molecular Compounds (cont’d) • First element in the compound is the regular element name • The last element in the molecule must have its name changed to an “ide” ending • The only time we are allowed to leave off a prefix is if there is only one atom of the first element in the molecule

47. Naming Molecular Compounds (cont’d) • Example: • C02 - carbon dioxide • CCl4 – carbon tetrachloride • P2O5 – diphosphorus pentaoxide

48. Your Turn • Complete sheet u3 s2 l2 Molecular Compounds

49. Ionic Bonding • An ionic bond is a force of attraction between a metallic ion and a non-metallic ion • It is an electron transfer between two elements or polyatomic ions • This type of bond created a crystal with a definite, repeated pattern

50. Consider that chemical reactions between metals and non-metals occur because metals lose their loosely held electrons to highly electronegative non-metals. • This results in ions of the opposite charge • These oppositely charged ions them mutually attract each other to form ionic compounds. • This attraction between oppositely charged ions is called an ionic bond.