Classification of Matter

1. Classification of Matter Matter Mixture Pure Substance Homogeneous (solution) Heterogeneous (mechanical mixture) Element Compound

2. Classification of Matter (Alternate) Matter Homogeneous Heterogeneous Mixture Pure Substance Element Compound

3. Physical properties • A physical property is any aspect of matter that can be measured or seen without changing the composition of the matter. • Qualitative • Odor, color, texture, state, malleability • Quantitative • Melting point, boiling point, density, mass

4. Physical Change • Does not change the composition of the matter - doesn't change what the substance is • In a mixture, individual components retain properties of the original mixture; e.g. dissolving salt in water • Change of state (melting, freezing, etc.) is an example of a physical change

5. Physical Separation • Separating a mixture based on physical properties • Involves a physical change only (not chemical) • Individual components have properties of the mixture • Examples: • Filtration • Using magnetic properties • Sedimentation, using density differences

6. Chemical Properties • A chemical property is any property of matter that becomes evident during a chemical reaction. • Can only be measured by changing a substance's chemical identity. • Chemical properties cannot be determined just by viewing or touching the matter.

7. Chemical Change • A new substance is formed and energy is either given off or absorbed • Involves energy • If heat is given off during the reaction, than the reaction is considered to be exothermic. • If heat is required for the reaction, than the reaction is considered to be endothermic. • Composition of the substance is altered • New substances are produced with properties different from the original substance • Not easily reversed

8. Evidence of Chemical Change • The following may indicate that a chemical change has occurred • Colour Change • Temperature Change • Odour Given Off • Precipitate is Formed • Gas Produced • Any new substance produced • Examples • Burning • Metal in acid • Electrolysis of water

9. Chemical Separation • Separating a substance using a chemical change • Can be used to separate compounds

10. The Atom • Three basic components of an atom: • Electrons • Protons • Neutrons • An atom is mostly empty space, with almost all of the mass contained in the nucleus • Protons and Neutrons are contained in the centre of the atom known as the nucleus. • Electrons “orbit” the nucleus.

11. Parts of the Atom • Neutron: Large with no charge (n0) • Proton: Large with a positive charge (p+) • Electron: Small with a negative charge (e-) • Charge on an electron is equal and opposite to the charge on a proton • All elements (in their ground state) are neutral, meaning the number of protons and the number of electrons are equal.

12. Representing Elements • To represent elements we use the symbol A X Z • X – Atomic symbol • Z – Atomic number (p+) • A – Mass number (p+ + n0)

13. Example • Na  Sodium • Atomic Number = 11  11 protons • Mass Number = 23  23 - 11 = 12 neutrons • Neutral  # electrons = # protons = 11 23 Na 11

14. Periodic Table • 118(?) Elements are arranged in Groups (columns) and Periods (rows) • There are three types of elements: • Metals • Metalloids • Non-Metals • Periodic Table is broken into sections for each type

15. Patterns and Trends Metals Metalloids Non-Metals Elements that border the “staircase” tend to have both metal and non-metal properties. These elements are known as metalloids.

16. Properties of the Types of Elements

17. Electron Configurations • Electrons “orbit” the nucleus in regions known as shells. • The elements in the first period have one shell and each period adds another. • The first shell can only hold 2 electrons and each shell after that can hold 8 electrons. • When a shell is full move on to the next one.

18. Examples of Electron Configurations

19. Outer Shell/Electrons • The outer most shell of an electron configuration is known as the valence shell. • Electrons contained in this shell are known as valence electrons. • What do you notice about elements of the same group and the number of electrons in the outer most shell?

20. Valence Electrons

21. Period Table and Configurations • Period determines the number of shells and the group determines the number of valence electrons. • When looking at chemical similarities in the periodic table we look at the groups and not the periods.

22. Common Names of Groups • Group 1 (1A) – Alkali Metals • Group 2 (2A) – Alkaline Earth Metals • Groups 3-12 are called the transition metals • Groups 1,2,13-18 are called representative elements

23. Common Names of Groups • Group 17 (7A) – Halogens • Means “salt former” • Group 18 (8A) – Noble Gases • Only for first six periods. • Called Noble Gases because they don’t easily react with the other elements (full valence shell)

24. Isotopes • Isotopes are variations of the element where neutrons are added or removed to give different types (weights) of atoms. • Still the same element just different mass number. • Atomic number does not change so there is still the same number of protons present. • The atomic mass seen on the periodic table is the average mass of all the isotopes of that element.

25. Isotope Example (Hydrogen) 0 neutrons 1 neutron 2 neutrons • 2H and 3H are known as deuterium and tritium. • Also called “heavy” hydrogen

26. Lewis Dot Diagrams • Lewis Dot Diagrams are used to just represent the valence electrons. • Hydrogen: • Helium: H He

27. How to Fill in Lewis Dot Diagrams • Start at the top and then fill in going clockwise. • Can only have up to a maximum of 8 dots around the atomic symbol. C N B

28. Charged Elements • Elements can become charged if the there is a change in the number of electrons • Elements try to get to the electron configuration of the closest noble gas –full valence shell. • Charged elements are known as ions. • Positively Charged Ion  Cation • Negatively Charged Ion  Anion • Boron has 3 electrons in its valence shell • Must lose 3 electrons to achieve a full valance shell • Acquires a charge of +3 • Ion is B3+ • How elements react depends on their valence electrons.

29. Positively Charged • Generally, metals tend to lose valence electrons relatively easily. • Elements that can easily lose an electron are known as electron donors. • To remove an electron from the Group 1 metals requires relatively little energy. • As you move down the group it becomes easier to remove the valence electron. • Metals are less reactive as you move across the period to the right.

30. Negatively Charged • Since non-metals have a greater number of valence electrons, they must gain electrons to fill their valence shell. • Elements that can easily gain an electron are known as electron acceptors. • The Halogens are very reactive elements. • As you move down the group the elements become less reactive. • Non-metals are less reactive as you move across the period to the left.

31. Forming Compounds • When two atoms collide, the valence electrons of each atom interact. • Elements try to get to the electron configuration of the closest noble gas - full valence shell.

32. Valence Shell • Three ways for an atom to acquire a full valence shell. • An atom may give up electrons • An atom may gain electrons • An atom may share electrons • When atoms give up and gain electrons in a reaction the resulting compound is known as an ionic compound with an ionic bond. • The third way to acquire a full valence shell will be talked about later in the course.

33. Ionic Compounds • Ionic compounds involve bonds between a metal cation and a non-metal anion. • If just two different elements are involved, than you have a binary compound. • Binary compounds require that the total charge (sum of the element’s charges) of the compound is equal to zero. • We represent the compound by writing down the element symbol for cation first and then the anion • Subscripts after each symbol identify how many ions are required for a total charge of zero. • The representation of the compound is known as the chemical formula

34. Lewis Structures • Sodium (Na+) and Chlorine (Cl-) • Now Chlorine has a full valence shell and the ionic compound NaCl is formed. Na Cl [ ]+ [ ]- Na Cl

35. Example • NaCl (Table Salt) • How do we name this compound? • Sodium Chloride • The suffix “ide” is put at the end of the name for the element that is the electron acceptor (anion) Na+ + Cl- NaCl • The sodium has a +1 charge and the chlorine has a -1 charge therefore +1 + -1 = 0.

36. Another Example • What would happen if we combined Magnesium and Chlorine? • Charges do not add up to zero. • Therefore we need more of one of the elements, but which one. • Magnesium has a 2+ charge and Chlorine has a 1- charge so we need two Chlorine. • MgCl2 (Magnesium Chloride) • Can also be done by drawing out the required number of atoms to get a total charge of zero.

37. Polyatomic Ions • Polyatomic Ions consist of two or more non-metal atoms grouped together. • There is only one common polyatomic cation • Ammonium NH4+ • There are several common polyatomic anions • Hydroxide OH- • Carbonate CO32- • Nitrate NO3- • Sulfate SO42- • Chlorate ClO3- • Phosphate PO43-

38. Polyatomic Ions • Compounds are named the same way • Writing the chemical formula is a little different - If more than one polyatomic ion is needed, than brackets must be put around the ion • Example: The chemical formula for ammonium oxide is (NH4)2O not NH42O • Do not forget the brackets!!!!

39. Polyatomic Example • Calcium Nitrate • Calcium (Ca2+) and Nitrate (NO3-) • Need two Nitrate ions to balance charges. • Ca(NO3)2

40. Transition Metals • Transition metals can form more than one ion - except for silver(+1), zinc (+2) and aluminum (+3). • For example Sodium can only produce the Na+ ion. Iron on the other hand can produce two ions. Fe  Fe2+ or Fe3+ • A roman numeral is placed after the atom in brackets to identify the charge • Iron that produces the +2 ion is iron(II) • Iron that produces the +3 ion is iron(III)

41. Examples • 1. Iron(III) Oxide (Rust) Fe3+ O2- Fe3+ O2- O2- • Charge of +6 from the iron and -6 from the oxygen. Chemical Formula - Fe2O3 • 2. CuCl2 Cu Cl- Cl- • Cu must have a +2 charge to balance the -2 from the 2 Cl. Copper(II) Chloride

42. Covalent Bonds • Two or more non-metallic elements. • Electrons must be shared since both atoms are looking to gain electrons. • When atoms share electrons they are joined by a covalent bond. • A neutral particle that is composed of atoms joined together by covalent bonds is called a molecule. • Substances that are composed of molecules are called molecular compounds.

43. Molecular Compounds • Water (H2O) • Two H+ atoms and a O2- atom. O H H O H H

44. Naming Molecular • H2O • Start with the element that is farther left on the periodic table (Hydrogen). • The rules for the second element still apply, suffix of “ide”. • Different is that the elements require prefixes depending on how many are in the compound. • So water’s chemical name is dihydrogen monoxide.

45. Prefixes • Prefix mono is only used for the second element. • “a” or “o” is left off of the prefix when used with an element starting with a vowel

46. Diatomic Molecules • Atoms can share electrons with the same atom. • These molecules have two of the same atoms joined by a covalent bond. • Since there are two of the same atoms the word diatomic is used. (“di” meaning two) • Seven elements exist as diatomics: • Hydrogen • Oxygen • Nitrogen • Fluorine • Chlorine • Bromine • Iodine

47. Ionic Compounds • Ionic compounds form large structures called lattices • Attraction between oppositely charged ions is strong.

48. Ionic Properties • Characteristics of an ionic compound: • Tend to have relatively high melting and boiling points because of the large amount of energy is needed to break the strong force of attraction in an ionic bond. • Conduct electricity when they are liquid or when they are dissolved in water. Melting or dissolving allow ions to move freely. In a solid state the ions are not able to move and therefore cannot conduct electricity.

49. Molecular Compounds • Bonds within the molecule are strong but forces of attraction between the molecules is weak.

50. Molecular Properties • Characteristics of a molecular compound: • Have relatively low melting points because little energy is needed to break the forces of attraction between molecules. • Relatively soft • Tend not to conduct electricity when they are in solid or liquid state. Do not conduct when dissolved in water because ions are not formed.