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Classification of Matter

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  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.