1. 1) Name that element! 1 2) Write formulas for: Calcium acetate Hydrogen peroxide Oxygen hexafluoride Hydroselenic acid

2. 1.(NH 4 ) 3 N 2.Cr(OH) 3 3.H 2 SO 3 4.NaClO 4 5.HCl 6.N 2 O 7.H 2 CO 3 8.SF 6 2

3. Protons: charge of +1 Neutrons: charge of 0 (neutral) Electrons: charge of -1 Neutral atoms have NO NET CHARGE Therefore, # protons = # electrons Subatomic Particles: Electric Charge 3

4. Subatomic Particles: Structure 4

5. Subatomic Particles: Mass Protons and neutrons have essentially the _______________. The mass of an electron is so small we ___________. Nucleus contains essentially _______________ of the atom. atomic mass unit (amu) – unit used to express masses of atoms; more convenient than grams (proper definition later) 1 g = 6.02214 x 10 23 amu 5

6. Symbols of Elements Elements are symbolized by one or two letters. 6

7. Atomic Number All atoms of the same element have the same number of protons so the number of protons defines that element. the atomic number = the number of protons 7

8. Mass Number mass number = number protons + number neutrons 8 of a specific atomin that atom

9. Isotopes Isotopes - atoms of the same element with different masses – The different masses are caused by different numbers of neutrons 11 6 C 12 6 C 13 6 C 14 6 C 9

10. Isotopes Compare and contrast Carbon-12 and Carbon-14 10 Carbon-12Carbon-14

11. Examples 11

12. Atomic Mass What’s the difference in atomic mass and mass number?? Elements exist in nature as a mixture of isotopes. – But wait! They each have a different mass, so which mass do we record on the periodic table? The atomic mass of each element is an average of the different isotope masses. Ex:- Hydrogen: 1.00794 amu - Oxygen: 15.9994 amu The amu is defined by assigning a mass of exactly 12 amu to the 12 C isotope of carbon. So 1 amu = ? – 1/12 th mass of single carbon-12 atom 12

13. At Haltom, we averaged grades just like in math. Add up your grades and divide by the number of grades. This is a typical average. Daily – 50% 63 Major – 50%+74 137 / 2 = 68.5 In this school district, they calculate grades like this: Daily grades – 40% 63 Major grades – 60% 74 Here since they aren’t equal weights, the average is more influenced by the major grades since they count as a higher %. This kid actually gets a 69.6  70 for the six weeks!

14. (% abundance x mass 1)+(% x mass 2)+(% x mass 3)+ etc.= average atomic mass for the element DO NOT DIVIDE BY ANYTHING! This is not a regular average. Don’t forget to change the percentage to a decimal first!!! means how often is it found

15. Gallium is an element found in lasers used in CD players. In a sample of gallium, there is 60.10% of 69 Ga (atomic mass 68.926) atoms and 39.90% of 71 Ga (atomic mass 70.925) atoms. What is the atomic mass of gallium? 15

16. An unknown element, Q, has an average atomic mass of 73.75 amu. The first isotope of Q has a mass of 72.99 amu and the second isotope has a mass of 74.99 amu. What are the percentages of each isotope?

17. 1s 2 2s 2 … 9/2/15 Write electron configurations and dot notations for the following elements: Silicon, Si and Cesium, Cs 17

18. Kinetics 9/3/15 Explain, in terms of molecular kinetics, why increasing the temperature of a reaction system always results in an increase in the rate of the reaction 18

19. How do we know about isotopes if we cant see them? 19 Steps: 1.Sample is vaporized. 2.Sample is ionized by knocking off 1 or more electrons (for easier manipulation by the electromagnet) 3.Charged (usually positive) particles are accelerated around a curve. 4.Electromagnet deflects particles (the lighter they are, the more they’re deflected; the more positive they are, the more they’re deflected). 5.Ions passing through the machine are detected. 6.Mass spectrum is produced.

20. Mass spectrum of which element? mass Boron

21. Mass spectrum of which element? mass Zinc

22. Periodic Table Elements arranged in order of increasing atomic number Rows = periods Columns = groups H 1 1.00794 atomic mass atomic symbol atomic number 22

23. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. 23

24. Groups Elements in the same group have similar chemical properties These five groups are known by their names. Memorize!! There are different numbering schemes for groups 24

25. Periodic Table Metals are on the left side of the chart. Examples of metallic properties: - luster - malleable - ductile - high electrical and heat conductivity - solids at room temp. (except Hg - liq) 25

26. Periodic Table Nonmetals are on the right side of the periodic table plus H. At room temp, some nonmetals are gaseous, some are are solids, and one is liquid (bromine). 26 Different from metals in appearance and other physical properties: dull, brittle, different colors, poor conductors of heat/electricity

27. Periodic Table Metalloids border the stair-step line (with the exception of Al, Po, and At). Metalloids have properties that lie between metals and nonmetals (semiconductors). 27

28. Examples a)Which 2 elements would you expect to show the greatest similarities in chemical and physical properties? S, Be, K, Ne, Se b)What is the symbol for arsenic? – Is it metal, nonmetal, or metalloid? – What group is it in? – What period? – What is its atomic number? – Atomic weight? – How many protons? 28

29. Atomic Mass 9/5/13 Rubidium has two common isotopes, 85 Rb and 87 Rb. If the abundance of 85 Rb is 72.2% (atomic mass 84.926) and the abundance of 87 Rb is 27.8% (atomic mass 87.113), what is the average atomic mass of rubidium? 72.2% 85 Rb, 27.8% 87 Rb Rubidium has two common isotopes, 85 Rb and 87 Rb. If the abundance of 85 Rb is 72.2% (atomic mass 84.926) and the abundance of 87 Rb is 27.8% (atomic mass 87.113), what is the average atomic mass of rubidium? 72.2% 85 Rb, 27.8% 87 Rb 29 (.722x 84.926) + (.278 X 87.113) = 85.5 amu

30. Molecules Molecule – an assembly of 2 or more nonmetals tightly bound together behaves as a single “package”; does not split apart when dissolved in water 30

31. Molecular Compounds Molecular compounds (or covalent compounds) are molecules that are made of more than one type of element 31

32. Diatomic Molecules Diatomic molecule – molecule made up of 2 atoms These seven elements occur naturally as molecules containing two atoms. “Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 ” – pronounced “Brinklehof” 32 Memorize!!

33. Types of Chemical Formulas Molecular formula gives the exact number of atoms of each element in a compound. Empirical formula gives the lowest whole-number ratio of atoms of each element in a compound. – Examples: Is H 2 O 2 a molecular or empirical formula? – Give the empirical formula for H 2 O 2 Give the empirical formula for glucose, whose formula is C 6 H 12 O 6. Give the empirical formula for carbon dioxide, CO 2. 33

34. Formulas: Example The structural formula of butane is What is the molecular formula of butane? What is the empirical formula of butane? 34

35. Ionic Compounds Ionic compound – a compound with both positively and negatively charged ions (e.g. NaCl) – Involves electron transfer between two neutral atoms – Generally formed between metals and nonmetals 35

36. Ions ion - a charged particle Forms when an atom loses or gains electron(s), resulting in a net charge Cations are positive ions – The “t” in cation looks like a + sign – Formed when an atom loses electrons (more protons than electrons, so +) Anions are negative ions – Formed when an atom gains electrons (more electrons than protons, so -) Nucleus does not take part in chemical reactions! Ions behave very differently from the neutral atom from which it was formed. Example: Na (sodium) 36

37. Ions: Example Chemical SymbolMass number# Protons# Neutrons# Electrons 75 Br - 1145254 1106346 Mg 2+ 20 37

38. 38 In which of the following groups are the three species isoelectronic; i.e., have the same number of electrons? a)S 2─, K +, Ca 2+ b)Sc, Ti, V 2+ c)O 2─, S 2─, Cl ─ d)Mg 2+, Ca 2+, Sr 2+ e)Cs, Ba 2+, La 3+ 2002 Multi-choice

39. Predicting Ionic Charge Atoms lose or gain electrons to obtain the same number of electrons as the nearest noble gas. 1+ 2+ 3+ 4+ or 4- 3-2-1- Generally, - Metals LOSE electrons to form CATIONS - Nonmetals GAIN electrons to form ANIONS 39

40. Predicting Ionic Charge: Examples Predict the most stable ion formed from the following elements: a)Potassium (K) b)Antimony (Sb) c)Oxygen (O) d)Aluminum (Al) Pg 55 & 56 in the book 40

41. Ionic and Molecular Compounds Predict whether the following would be considered ionic or molecular compounds. a)MgCl 2 b)CO 2 c)SF 6 d)Ca 3 N 2 41

42. Chemical Nomenclature (Naming Compounds) Systematic set of rules Inorganic compounds a) ionic compounds b) binary molecular (or covalent) compounds c) acids 42

43. Writing Formulas: Ionic Compounds 1.Make sure it is ionic by identifying cation and anion 2.Write the ________ first and then _________, with charges. 3.Use “criss-cross” method to write subscripts The charge on the cation becomes the subscript on the anion The charge on the anion becomes the subscript on the cation Do not write subscripts of 1 – they are understood If either ion is polyatomic AND NEEDS A SUBSCRIPT GREATER THAN 1, place parentheses around the polyatomic with the subscript outside the parentheses – If the subscript for the polyatomic is 1, then neither the parentheses nor the subscript is necessary 4.If the new subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor 43

44. Compounds 9/6/13 Write the chemical formula for the following 1.Chlorine gas 2.Ammonium Sulfate 3.Chromium (II) Sulfide Name the following 1.Fe 2 O 3 2.MgSO 3 3.Na 3 PO 4 44

45. Naming Ionic Compounds: Cations 1.Is the cation…? a)NH 4 + ? Just write ammonium! b)metal ? Pick either choice i. or ii. i.A metal ion that can have only one possible charge: -Just write the metal name! Li + lithium Al 3+ aluminum ii. A metal ion that can have more than one possible charge: -Write the metal name followed by the charge indicated by a Roman numeral in parentheses -Write Latin name with ending to show charge (-ous = lower charge; -ic = higher charge) -These are often transition metals Cu + copper(I) ion or cuprous ionFe 2+ iron(II) ion or ferrous ion Cu 2+ copper(II) ion or cupric ionFe 3+ iron(III) ion ferric ion 45 OR

46. Naming Ionic Compounds: Anions 2.Is the anion…? a) polyatomic ? Memorize common ion list on last page of book Write the polyatomic name without changing it b)nonmetal (monatomic) ? Write the name of the nonmetal, replacing the ending with -ide: H - hydride O 2- oxide N 3- nitride 46

47. Summary to Naming Ionic Compounds: Writing the Name from the Formula 1.Make sure it’s ionic!!! 2.Identify the cation and anion. 3.Write the name of the cation. – If the cation is an metal ion that can only have one possible charge, write the name of the metal – If the cation can have more than one possible charge, write the name of the metal, followed by the charge as a Roman numeral in parentheses – If the cation is NH 4 +, write ammonium 4.Write the name of the anion. – If the anion is an element, change its ending to -ide – If the anion is a polyatomic ion, simply write the name of the polyatomic ion 47 Does the metal have only 1 possible charge or more than 1? Group 1A, 2A, 3A, Zn 2+, Ag +, Cd 2+, or NH 4 + Other transition metals, Groups 4A, 5A 1.Write the metal name or ammonium for the NH 4 + ion. 1. Write the metal name followed by a Roman numeral in parentheses for the positive charge of the ion. Is the anion formed from a single atom or a polyatomic? 2. Write the element name, changing the ending to –ide. 2. Write the name of the polyatomic ion. AnionMetal (or NH 4 + )

48. Examples: Write the name a)Fe 2 O 3 b)(NH 4 ) 2 CO 3 c)CuCl 2 d)Zn 3 P 2 e)Cr 2 (SO 3 ) 3 f)K 2 S g)SnO 2 48

49. Examples: Write the formula a)Cu 2+ and O 2- b)Al 3+ and S 2- c)Ca 2+ and ClO 3 - d)iron(III) hydride e)sodium sulfate f)stannous nitride What was once believed to be "Iron(III) hydride" or "ferric hydride"(FeH 3 ) was later shown to be FeH bound to molecular hydrogen H 2. 49

50. Naming9/9/13 Write the names or formulas 1.Cr(OH) 3 2.Fe(NO 3 ) 2 3.zinc acetate tetrahydrate 4.potassium dichromate Write the names or formulas 1.Cr(OH) 3 2.Fe(NO 3 ) 2 3.zinc acetate tetrahydrate 4.potassium dichromate 50 chromium (III) hydroxide or chromic hydroxide iron (II) nitrate or ferrous nitrate Zn(C 2 H 3 O 2 ) 2 4H 2 O K 2 Cr 2 O 7

51. Naming Acids Acid – substance whose molecules yield hydrogen ions (H + ) when dissolved in water -In the form HX, where X = an anion To name acids, first determine name of anion (3 possibilities): If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- HClHBrHI If the anion in the acid ends in -ite, change the ending to -ous acid HClO (hypochlorite ion)HClO 2 (chlorite ion) If the anion in the acid ends in -ate, change the ending to -ic acid HClO 3 HClO 4 (perchlorate ion) 51

52. 1)The name of the element farther to the left in the periodic table is usually written first. EXCEPT: Oxygen is always written last, unless fluorine is included – then F is last Example: CO 2 or PCl 3 2)If both elements are in the same group, the one having the higher atomic number is named first. Example: ICl 3 Binary molecular compounds – composed of 2 nonmetals 52 Naming Binary Molecular Compounds

53. 3)Change name of second element to -ide Ex: carbon monoxide 4)Add prefix to denote the number of atoms of each element - mono- is never used on the first element listed - When adding the prefix results in an a and/or o together, drop the first vowel Ex: PF 5 phosphorus pentafluoride Cl 2 O dichlorine monoxide N 2 O 4 dinitrogen tetroxide 53 Naming Binary Molecular Compounds

54. Write names for the following compounds: 1.CaO 2.(NH 4 ) 3 N 3.Cr(OH) 3 4.H 2 SO 3 5.NaClO 4 6.HCl 7.N 2 O 8.H 2 CO 3 9.SF 6 10.Cl 2 O 7 54 Examples Write formulas for the following compounds: 1.iron(III) chloride 2.disulfur dichloride 3.hydrosulfuric acid 4.cobalt(II) phosphate 5. tetraphosphorus decoxide