Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Matter All things are composed of matter. Kinetic Theory of Matter – All matter is composed of tiny particles that are in constant motion. Phases of matter 1. Solid a. definite shape and volume b. can be crystallized c. closely packed slow moving particles 2. Liquid a. definite volume but takes the shape of the container b. closely packed particles that flow past each other c. cannot be compressed 3. Gas a. takes the shape and volume of its container b. can be compressed c. fast moving particles separated by space 4. Plasma a. fastest moving particles b. high temperature ionized gas c. composes 90% of the matter in the universe Changes of phase 1. evaporation – liquid to a gas below boiling point 2. condensation – gas to a liquid 3. boiling – liquid to a gas at boiling point 4. melting – solid to liquid at melting point 5. freezing – liquid to solid at freezing point 6. sublimation – solid to gas or reverse 7. Each of these changes of phase results in the gain or loss of energy Physical properties – Can be observed without changing the substance a. melting/freezing point b. boiling point c. color d. odor e. ductility f. malleability g. solubility h. conductivity Physical changes – Does not change the composition to the matter a. change in state b. change in shape c. change in size Chemical properties – Seen when the substance changes into another substance a. reactivity b. combustibility Chemical change – Change that produces a new substance. Recognized when the following occurs a. gas production b. formation of a precipitate c. burning d. rusting Classification of matter 1. substance – same throughout, cannot be separated by physical means a. compound – more than one type of atoms b. element – only one type of atom 2. mixture – not the same throughout, can be separated by physical means a. homogeneous – same throughout 1. solutions – particles evenly distributed 2. Ex: salt water, iced tea b. heterogeneous – not the same throughout 1. suspension – settles out Ex: muddy water 2. colloid – never settles out Ex: milk, smoke Methods of Separating Mixtures 1. physically – using hand to sort 2. filtration – using a membrane (filter) to catch large particles 3. magnetism – using a magnet to pick-up materials 4. evaporation – solid remains after liquid evaporates 5. fractional distillation – heating a mixture of liquids with different boiling points and collecting the vapors released at each boiling point Atomic Theory Scientist A. 400 B.C. – Democritus B. Aristotle C. early 1800’s – John Dalton D. 1897 – J.J. Thompson E. 1912-1913 - Lord Ernest Rutherford John Dalton (1807) Billiard Ball Model atom is a uniform, solid sphere J. J. Thomson (1903) Cathode Ray Tube Experiments Discovered Electrons beam of negative particles negative particles within the atom Plum-pudding Model J. J. Thomson (1903) Plum-pudding Model positive sphere (pudding) with negative electrons (plums) dispersed throughout Ernest Rutherford (1911) Gold Foil Experiment Discovered the nucleus dense, positive charge in the center of the atom Nuclear Model Ernest Rutherford (1911) Nuclear Model dense, positive nucleus surrounded by negative electrons Niels Bohr (1913) Bright-Line Spectrum Energy Levels tried to explain presence of specific colors in hydrogen’s spectrum electrons can only exist in specific energy states Planetary Model Niels Bohr (1913) Bright-line spectrum Planetary Model electrons move in circular orbits within specific energy levels 1912-1913 - Lord Ernest Rutherford 1. The Gold Experiment – A sheet of gold foil was bombarded with positively charged subatomic particles. a. Most particle passed through. b. A few particle were deflected at large angles. c. About 1 in 8000 bounced back in almost the opposite direction. Rutherford’s Experiment http://glencoe.mcgrawhill.com/sites/0078807220/student_view0/ chapter28/concepts_in_motion.html Elements A. Have their own characteristics 1. physical properties 2. chemical properties B. Unique name and symbol 1. universally accepted 2. symbol is 1-3 letter abbreviation a. 1st letter is always capitalized b. 2nd and 3rd letters are always lower case Make-up of atoms 1. nucleus a. protons – positive with 1 amu of mass – ID’s element b. neutron – no charge with 1 amu of mass – ID’s isotope 2. electron clouds a. area where you are most likely to find electrons b. electrons – negative particles with extremely small mass – ID’s ion D. Atomic number (Z) – number of protons E. Atomic mass (A) – number of protons plus the number of neutrons # of neutrons = A – Z F. Electrically neutral atoms have equal numbers of protons and electrons G. Ions 1. positive - # of proton > # of electrons 2. negative - # of electron > # of protons H. Isotopes – atoms of the same element with different numbers of neutrons. History of the periodic table 1869 Dimitri Mendeleev – Russian chemist, proposed the use of atomic weight to classify elements. His table was more widely accepted than Meyer’s table. He used the properties of the elements to group them. He had blank spaces, that he predicted what the element would be like. E. 1913 Henry Mosely – English scientist, proposed the use of atomic numbers to classify elements as a modification of Mendeleev’s Table. This cleared up inconsistencies in Mendeleev’s Table. Using the Periodic Table A. groups or families are in columns (18) B. periods are in rows (7) C. stair steps separate metals and nonmetals Metals 1. to the left and below the stair steps 2. properties a. shiny surface b. good conductors c. give up electrons to form compounds d. usually have 3 or less electrons in outer shell. e. can be ductile, malleable or brittle Metal Families a. alkali metals – Group 1 except hydrogen 1. most reactive 2. exist only as compounds in nature 3. good conductors, shiny, ductile, and malleable 4. filling s orbital b. alkaline earth metals – Group 2 1. second most reactive group of metals 2. exist only as compounds in nature 3. filling s orbital c. transition metals – Groups 3-12 1. most commonly known metals 2. contains the best conductors: Copper, Silver, and Gold 3. contains the natural magnets: Iron, Nickel, and Cobalt 4. contains the most unreactive metals: copper, silver, and gold 5. contains a liquid metal: mercury 6. filling d orbitals d. rare earth metals 1. 2 series i. lanthanide ii. actinide 2. filling f orbitals 4. Reactivity for metals increases form right to left and top to bottom. Nonmetals 1. to the right and above the stair steps 2. Properties a. dull surface b. poor conductors c. share or gain electrons to form compounds d. usually have 3 or more electrons in outer shell 3. families a. hydrogen i. diatomic molecule in nature (H2) ii. highly reactive iii. can gain or lose electrons b. halogens – Group 17 i. most reactive of nonmetals – needs one electron to be stable ii. range from gases to liquids to solids c. noble gases – Group 18 i. most stable – outer orbital is filled ii. gases at room temperature iii. originally called inert gases because they were thought not to react Mixed Families a. Boron family – Group 13 i. one metalloid – boron (B) ii. four metals iii. three metals used for semiconductors – Ga, In, and Tl iv. contains the most abundant element in the earth’s crust – Al b. Carbon family – Group 14 i. one nonmetal – carbon (C) ii. two metalloids – Si and Ge iii. two metals – Sn and Pb c. Nitrogen family – Group 15 i. two nonmetals – N and P ii. two metalloids – As and Sb iii. one metal – Bi d. Oxygen family – Group 16 i. three nonmetals – O, S, Se ii. two metalloids – Te and Po 5. Reactivity increases from left to right and bottom to top. Metalloids 1. along the stair steps 2. properties of both metals and nonmetals 3. There are eight – B, Si, Ge, As, Sb, Po, Te, and At 4. Boron is considered the least reactive nonmetal because it has three electrons in its outer shell and needs five more to be stable Chemical Bonding A. Metallic bonding – metal atoms closely attracted to each other allow their electrons to flow from one atom to another. B. Ionic bonding 1. Ion – atom or group of atoms with a charge ( + or - ) 2. Attraction between ions of different charges (+ or - ) 3. Weak bond that is easily broken. 4. Can be form by the exchange of electrons a. one atom gains electrons b. one atom loses electrons C. Covalent Bonding 1. Strong bond 2. Sharing of electrons by two atoms. Writing and naming compounds A. Naming binary compound 1. Write the name of the first element or polyatomic ion. 2. Write the name of the second element changing the ending to –ide, or write the name of the polyatomic ion. B. Writing formulas 1. Cancel change method a. Write the element symbol for the first element or polyatomic ion b. Write the element symbol for the second element or polyatomic ion c. Use subscripts to indicate how many of each type of ion is needed to cancel the charges d. Use parentheses if more than one of a polyatomic ion is needed. EX: sodium chloride Na Cl + +1 + (-1) = 0 Charges cancelled Formula is NaCl EX: aluminum chloride Al Cl +3 +3 + (-1) = 2 Charges are not cancelled. +3 + (3) (-1) = 0 Charges are cancelled. Formula is AlCl3. 2. Cross over method a. Write the formula for each ion b. If the charges are opposite but equal in value, drop the charges and write the formulas together. EX: lithium chloride Li+ ClFormula LiCl c. If the charges are not equal in value, drop the sign for the charge and use the number from the charge as the subscript for the opposite ion. EX: aluminum chloride Al+3 Cl↓←←←←↑ Al3 Cl ↓→→→→↑ AlCl3 formula d. If a metal has a Roman numerals in parentheses after its name, the metal can form more than one ion and the Roman numeral indicates the metals positive charge. EX: iron (III) sulfide Fe+3 S-2 ↓←←←↑ Fe3 S2 ↓→→→↑ Fe2S3 formula e. When polyatomic ions are part of the formula, use parenthesis if there is more than one polyatomic ion. EX: mercury (II) phosphate Hg+2 PO4-3 ↓←←←←↑ Hg 2 PO4 3 ↓→→→→↑ Hg 3 PO4 2 Hg3(PO4)2 formula EX: hydronium phosphate H3O+ PO4-3 ↓→→→↑ H3O PO4 3 ↓→→→↑ H3O 3 PO4 (H3O)3PO4 ← formula C. Naming compounds containing metals with more than one oxidation state. 1. Oxidation state refers to the charge for the ion (the number of electrons that can be gained or lost) 2. Determine the charge of the negative ion. EX: Fe2O3 Oxygen is always -2 3. Multiply the negative charge by the subscript on the negative ion. EX: Fe2O3 Charge is -2. Subscript is 3. -2 x 3 = -6 4. Drop the negative sign and divide by the subscript of the positive ion. EX: Fe2O3 Total negative charge is -6. Positive subscript is 2 6 ÷ 2 = 3 5. Write the name of the metal and in parentheses using a Roman numeral write the positive charge (oxidation state) just calculated. EX: Fe2O3 iron (III) 6. Then write the name of the negative ion EX: Fe2O3 iron (III) oxide 7. Roman numerals 1 I 6 VI 2 II 7 VII 3 III 8 VIII 4 IV 9 IX 5 V 10 X 8. Naming negative ions a. Drop the ending on the element name and add –ide. b. hydrogen nitrogen oxygen sulfur carbon phosphorus HN-3 O-2 S-2 C-4 P-3 hydride nitride oxide sulfide carbide phosphide chlorine fluorine iodine bromine arsenic silicon selenium astatine tellurium Cl FIBr As -3 Si -4 Se -2 At Te-2 chloride fluoride iodide bromide arsenide silicide selenide astatide telluride VIII. Lewis Structures – Electron Dot Diagrams A. Show the placement of valance electrons B. Used to explain bonding angles and crystal structures C. Writing electron dot diagrams 1. Determine the number of valance electrons the element has a. Group 1 has 1 b. Group 2 has 2 c. Group 13 has 3 d. Group 14 has 4 e. Group 15 has 5 f. Group 16 has 6 g. Group 17 has 7 h. Group 18 has 8, except He which has 2 i. For the transition metals and rare earth metals refer to their electron configuration. The number of valance electrons is always the last number. 2. Write the element symbol 3. Start on one side of the element symbol, place one dot on each side before pairing to equal the total number of valance electrons. EX: hydrogen Group 1 element → has 1 valance electron EX: phosphorus Group 15 element → has 5 valance electrons IX. Formula mass – The sum of the masses of all atoms making up a molecule. A. Calculating formula mass 1. Determine what elements are present a. determine how many atoms of each element are present b. look up the average atomic mass for the element and round to the nearest .5 amu. 2. Calculate the total mass of each element present by multiplying the number of atoms of an element by the element’s average atomic mass. 3. Calculate the total mass of the molecule by adding the total masses of all the elements in the molecule. B. Examples 1. FeO Element # of atoms Average atomic mass Total mass for element Formula Mass Fe 1 O 1 55.847 15.999 55.847 71.846 15.999 (Fe)(1) + (O)(1) = FeO (55.847) (1) + (15.999)(1) = 71.846 Element # of atoms Fe 1 x O 1 x Formula mass mass t. mass 55.847 = 55.847 15.999 = +15.999 = 71.846 2. NaOH Element # of atoms Average atomic mass Total mass for element Formula mass Na 1 O 1 H 1 22.990 15.999 1.0079 22.990 39.9969 15.999 1.0079 (Na)(1) + (O)(1) + (H)(1) = NaOH (22.990)(1) + (15.999)(1) + (1.0079)(1) = 39.9969 Element # of atoms Na 1 x O 1 x H 1 x Formula mass mass t. mass 22.99 =22.99 15.999 = 15.999 1.0079= +1.0079 = 39.9969 3. AlPO4 Element Al Number of atoms 1 Average atomic mass 26.983 Total mass of element 26.982 Formula mass 121.952 P 1 O 4 30.974 15.999 30.974 63.996 (Al)(1) + (P)(1) + (O)(4) = AlPO4 (26.983)(1) + (30.974)(1) + (15.999)(4) = 121.952 Element # of atoms Al 1 x P 1 x O 4 x Formula mass mass t. mass 26.983 = 26.983 30.974 = 30.974 15.999 = +63.996 = 121.952 4. Hg3(PO4)2 Element Number of atoms Average atomic mass Total mass of element Formula mass Hg 3 P 2 O 8 200.59 30.974 15.999 601.77 61.948 127.992 791.71 (Hg)(3) + (P)(2) + (O)(8) = Hg3(PO4)2 (200.59)(3) + (30.974)(2) + (15.999)(8) = 791.71 Element # of Hg P O Formula mass atoms 3 2 8 mass t.mass 200.59 = 601.77 30.974 = 61.948 15.999 = 127.992 791.71