1. CHAPTER 1 NOTES ATOMS AND BONDING

2. Section 1: Elements and Atoms 1. Elements and Atoms A. Matter is defined as anything that has mass and takes up space. - Ancient Greeks believed that all matter was made up of four elements; earth, air, fire and water.

3. B.Elements are the simplest pure substances, and they cannot be broken down into any other substance. (the periodic table of elements) - Often called the building blocks of matter because all matter is composed of one element or a combination of two or more elements. C. Compound – a pure substance made of two or more elements that are chemically combined. Ex. Sodium Chloride (table salt) = NaCl

4. D. Mixture – two or more substances that are in the same place but are not chemically combined. Ex. Air, soil, wood, orange juice

5. E. Atoms 1. Democritus (430 B.C.) Proposed the idea of an atom. 2. Matter is formed of pieces that are “uncuttable” 3. Modern definition: a. An atom is the smallest particle of an element.

6. 2. Atomic Theory and Models A scientific theory- a well-tested idea that explains and connects a wide range of observations Daltons atomic theory- 1. All elements are composed of atoms. 2. All atoms of the same element are exactly alike and have the same mass. Atoms of different elements are different and have different masses. 3. An atom of one element cannot be changed into an atom of a different element. Atoms cannot be created or destroyed in any chemical change, only rearranged. 4. Every compound is composed of atoms of different elements, combined in a specific ratio.

7. Thomson- 1. Found atoms contained negatively charged particles called electrons. Aka, “muffin model”

8. Rutherford- gold foil experiment 1. Inferred that an atoms positive charge must be clustered in a tiny region in its center called the nucleus 2. The nucleus is made of one or more positively charged particles called protons

9. Bohr’s Model-electrons orbit the nucleus in certain orbits similar to the way planets orbit the sun.

10. Electron Cloud Model-electrons can be anywhere in a cloudlike region around the nucleus. The “cloud” symbolizes where an electron is likely to be found. An electron’s movement is related to its energy level- specific amount of energy it has. Electrons with more energy are farther from the nucleus. Modern Atomic Model-James Chadwick discovered the neutron, a neutral particle in the nucleus with nearly the same mass as a proton.

11. SOME QUICK FACTS….  The mass of an atom is extremely small. The units of mass used to describe atomic particles is the atomic mass unit (or amu ).  An atomic mass unit is equal to 1.66054 x 10 -24 grams.  How do the different sub-atomic particles compare as far as their mass?  Proton = 1.0073 amu  Neutron = 1.0087 amu  Electron = 5.486 x 10 -4 amu  From this comparison we can see that: The mass of the proton and neutron are nearly identical  The nucleus (protons plus neutrons) contains virtually all of the mass of the atom  The electrons, while equal and opposite in charge to the protons, have only 0.05% the mass  The size of an atom is quite small also, the typical range for atomic diameters is between 1 x 10 -10 and 5 x 10 -10 meters.  Note: a convenient unit of measurement for atomic distances is the angstrom (Å). The angstrom is equal to 1 x 10 -10 meters. Thus, most atoms are between 1 and 5 angstroms in diameter.

12. SECTION 2: ATOMS, BONDING, AND THE PERIODIC TABLE

13. 1. Valence electrons and Bonding a. valence electrons- electrons that are in the highest energy level of an atom and are held loosely by the nucleus -valence electrons determine the properties of the element and how the atom can bond with other elements

14. b. electron dot diagrams-symbol of an element surrounded by the number of dots that equals the number of valence electrons the element has.

15. c. Chemical bond- the force of attraction that holds two atoms together as a result of the rearrangement of electrons between them -atoms are less likely to react with they have 8 valence electrons. -atoms will gain or lose electrons to become more stable. -when atoms bond, electrons can be transferred from one atom to another or are shared. -the bonding of atoms causes a chemical reaction and makes a new substance.

16. 2. The Periodic Table A. First put together in the 1800’s by a Russian Chemist –Dimitri Mendeleev B. Modern periodic table is used worldwide 1. Elements are organized into groups with the same number of valence electrons in their atoms. 2. Important information contained within each block

18. atomic number = number of protons 1) protons are positively charged subatomic particles 2) nucleus is found in the center of the atom. It is made up of protons and neutrons. 3) neutron – is a neutral particle – nearly the same mass as a proton. 4) electron – typically equals the number of protons in the Atom, except if the atom is an ion, it is a negative particle that orbits around the nucleus in different energy levels. 5). atomic mass = the number of protons and neutrons in the nucleus of an atom.

19. Calculations involving Subatomic Particles: atomic number = # of protons mass number = # of protons + # of neutrons neutral atom: # of protons = # of electrons charged ion: # of electrons = # of protons - +charge # of electrons = # of protons + - charge

20. 3. Organization of the periodic table (see periodic table handout) Groups (family)– a column (vertical) of elements that have the same number of valence electrons. Period – elements in the same horizontal row of the periodic table. Elements are classified into three major categories that include metals, nonmetals, and metalloids.

21. 4. Families of the periodic table a. alkali metals-very reactive; Li, Na, K, Rb, Cs, Fr b. alkaline earth- reactive; Be, Mg, Ca, Sr, Ba, Ra c. halogens-“salt formers”; F, Cl, Br, I d. noble gases- all stable, few react; He, NE, Ar, Kr, Xe, Rn

22. SECTION 3 IONIC BONDS

23. REVIEW….

24. 1. Ions and ionic bonds a. ion- an atom or group of atoms with a charge. -cation- an atom that loses an electrons and has a positive charge. -anion- an atom that gains an electrons and has a negative charge. b. polyatomic ion- ions that are made of more than one atom and can be positive or negative. c. ionic bonds-attraction between oppositely charged ions. d. an ionic compound is one that consists of positive and negative ions. e. cations are usually metals and anions are usually nonmetals

25. 2. Chemical formulas and names A. Formula writing Ex: sodium chloride  NaCl Magnesium chloride  MgCl 2

26. Group 11+Group 162- Group 22+Group 171- Group 153-Group 180 Formula Unit-Simplest ratio of the ions represented in an ionic compound. Formula units have a net zero charge Determining Charge 1.monoatomic ions –a one atom ion 2.Charge based on the location of the element on the periodic table 3. oxidation number –charge of the monatomic ion

27. OXIDATION NUMBERS

28. DETERMINING THE FORMULA : Criss-cross method 1. Cations listed first and anions listed second 2. Find the symbols and charges in the upper right corner of the symbol 3. Cross the charges, drop the (+) or (-) sign 4. Reduce if necessary and never write the number 1

29. Examples:  potassium oxide ______________________  aluminum sulfide ______________________  calcium fluoride _______________________  zinc fluoride ________________________  aluminum iodide _______________________

30. Examples - Answers  potassium oxide _______ K 2 O_ ________  aluminum sulfide __________ Al 2 S 3 ____________  calcium fluoride __________ CaF 2 _____________  zinc fluoride _____________ ZnF 2 ___________  aluminum iodide _____________ AlI 3 __________

31. NAMING BINARY IONIC COMPOUNDS  1. The name of the positive ion always come first.  2. The name of the positive ion (metal) comes directly from the periodic table.  3. The name of the negative ion (non-metal) is made by taking the root of the elements name and the suffix “ide”.  Examples:  Write the name of the following formula: H 2 S  1. Write the name of the positive ion first =  hydrogen  2. Write the root of the second element =  “sulf”- plus “ide” = sulfide  3. Name = Hydrogen sulfide 

32. PRACTICE  Write the name of the following formula: MgBr 2  1. positive ion: __________________   2. root of second element plus “ide”: ___________________   3. compound name: ______________________________  Magnesium bromide

33. IONIC COMPOUNDS CONTAINING POLYATOMIC IONS (PAI’S)  A. Polyatomic Ions  1. certain ions made up of covalently bonded  atoms that tend to stay together.  2. PAI’s act like a single atom when combining  with other atoms.  3. although the bonds of the PAI are covalent,  they tend to form ionic bonds with other  atoms.

34. NAMING SIMPLE IONIC COMPOUNDS CONTAINING PAI’S  Rules  a. Name the metal (positive ion) first. Use the name directly from the periodic table.  b. Name the PAI. Use your reference sheet for the proper name.  c. Write the metal first. The PAI is written second.  Examples:  NaOH = ______________________  Ca(NO 3 ) 2 = ____________________

35. ANSWERS  NaOH = sodium hydroxide  Ca(NO 3 ) 2 = calcium nitrate

36. REMINDER - FOR COMPOUNDS CONTAINING PAI’S II. Polyatomic ions each have specific names which must be recognized so you can use your chart. (At this point, if you are asked to name any compound that contains more than two elements, it will contain at least one polyatomic ion.) USE YOUR CHART! FormulaName C 2 H 3 O 2 1- acetate CO 3 2- carbonate HCO 3 1- bicarbonate NH 4 1+ ammonium A few of the more common polyatomic ions FormulaName NO 3 1- nitrate OH 1- hydroxide PO 4 3- phosphate SO 4 2- sulfate

37. FORMING IONIC COMPOUNDS WITH PAI’S  Rules  a. identify what type of atoms/PAI’s are in the compound.  b. write each atoms / PAI’s oxidation number.  c. determine how many of each atom/PAI is needed to make the compound neutral.  ***If more than one unit of the PAI is needed, be sure to enclose the PAI in parenthesis and write the correct subscript.  Examples:  boron phosphide = BPO 3  barium permanganate = Ba(MnO 4 ) 2 

38. FORMULA WRITING FOR POLYATOMIC IONS (SEE HANDOUT FOR LIST) 1. Write the element symbols and charges 2. Cross the charges and drop the sign 3. If the oxidation number going to the polyatomic is greater than 1, add parenthesis (keep the team together!). 4. Reduce the charges if possible, but do not change the subscripts on the polyatomic ion.

39. EXAMPLES OF FORMULA WRITING FOR PAI’S  sodium nitrate ______________________  potassium acetate ___________________  aluminum sulfate_____________________  lithium chlorate ______________________

40.  sodium nitrate __________ NaNO 3 _________  potassium acetate _______ KC 2 H 3 O 2 ________  aluminum sulfate_______ Al 2 (SO 4 ) 3 _________  lithium chlorate ________ LiClO 3 ___________  POLYATOMIC IONS!! (SEE HANDOUT FOR CHARGE)

41. NAMING IONIC COMPOUNDS CONTAINING A TRANSITION METAL  Rules:  1. The name of the positive ion is always written first.  2. The name of the positive ion is taken directly from the periodic table.  3. The name of the positive ion is followed by Roman numeral written in  parenthesis which equals the charge of the ion. There is no space between  the name and the parenthesis.  4. The name of the negative ion is made by taking the root of the elements  name and adding the suffix “ide”.

42. EXAMPLE: Write the name of the formula: CuCl 2  1. Write the name of the positive ion: copper  2. Determine the charge of the copper ion based on the charge of the nonmetal involved in the compound.  In this case, two chlorides have a charge of –1 each, a total of –2. In order for the compound to balance to zero, the charge on copper must be +2, so the Roman numeral for the charge is (II).  3. Write the root of the second element: “chlor” plus “ide” = chloride  4. Name = copper(II) chloride  Write the name for  Fe 2 O 3 =  Iron (III) oxide  PbO =  Lead (II) oxide

43. TRANSITION METALS – HELPFUL HINTS FOR NAMING. 1)Name the cation(metal) first. 2)Monoatomic cations (metals) use the element name 3)Monoatomic anions (non-metals)take their name from the root of the element name and add –ide 4)Some metals can have multiple oxidation numbers – distinguish them by using Roman numerals  Ex. Fe+2 and Fe+3 iron (II)iron (III) 5)Group I metals always +1 and Group II metals always +2 6)If the compound contains a polyatomic ion, simply name the Polyatomic ion (use your chart).

44. FORMULA WRITING FOR IONIC COMPOUNDS WITH TRANSITION METALS If a Roman numeral is required, the charge on the metal ion must be determined from the charge on the negative ion. Helpful Rules to Remember A metal ion is always positive. The Roman numeral indicates the charge, not the subscript. The positive and negative charges must cancel (total charge must = 0). Nonmetals are always negative & can never form more than one monatomic ion.

45. Examples

46. PRACTICE EXAMPLES Be(OH) 2 Cu 2 S SrS ZnI 2  Beryllium hydroxide  Copper (I) Sulfide  Strontium sulfide  Zinc (II) iodide

47. FORMULA WRITING WITH TRANSITION METALS Iron (III) phosphide = FeP Chromium (III) oxide = Cr 2 O 3 Nickel (II) chloride = NiCl 2

48. PRACTICE…

49. ADDITIONAL EXAMPLES: PAI AND TRANSITION METALS  Na 2 SO 4 _________________________  Fe(NO 3 ) 2 _________________________  AlCl 3 _________________________  PbI 4 _________________________  (NH 4 ) 3 PO 4 _________________________  Mg 3 N 2 _________________________ C 2 H 3 O 2 1- acetate CO 3 2- carbonate HCO 3 1- bicarbonate NH 4 1+ ammonium NO 3 1- nitrate OH 1- hydroxide PO 4 3- phosphate SO 4 2- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

50.  AgC 2 H 3 O 2 ________________  NaOH ___________________  NaClO 3 ___________________  NaC 2 H 3 O 2 _________________  BaSO 4 ___________________  ZnCO 3 ___________________ C 2 H 3 O 2 1- acetate CO 3 2- carbonate HCO 3 1- bicarbonate NH 4 1+ ammonium NO 3 1- nitrate OH 1- hydroxide PO 4 3- phosphate SO 4 2- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

51. ADDITIONAL EXAMPLES: PAI AND TRANSITION METALS Na 2 SO 4 sodium sulfate Fe(NO 3 ) 2 iron (II) nitrate AlCl 3 aluminum chloride PbI 4 lead (IV) iodide (NH 4 ) 3 PO 4 ammonium phosphate Mg 3 N 2 magnesium nitride C 2 H 3 O 2 1- acetate CO 3 2- carbonate HCO 3 1- bicarbonate NH 4 1+ ammonium NO 3 1- nitrate OH 1- hydroxide PO 4 3- phosphate SO 4 2- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral. AgC 2 H 3 O 2 Silver (I)acetate

52.  NaOH Sodium hydroxide  NaClO 3 Sodium chlorate  NaC 2 H 3 O 2 Sodium acetate  BaSO 4 Barium sulfate  ZnCO 3 zinc (II) carbonate C 2 H 3 O 2 1- acetate CO 3 2- carbonate HCO 3 1- bicarbonate NH 4 1+ ammonium NO 3 1- nitrate OH 1- hydroxide PO 4 3- phosphate SO 4 2- sulfate * Groups I & II, Al, Zn, Cd, and Ag need no Roman numeral.

53. PROPERTIES OF IONIC COMPOUNDS a. hard, brittle crystals b. have high melting points c. solids at room temperature d. conduct electricity when dissolved in water

54. SECTION 4 COVALENT BONDS

55. 1. How Covalent bonds form a. When two atoms share electrons it is called a covalent bond b. Covalent bonds form between nonmetals c. A molecule is a neutral group of atoms formed by covalent bonds d. Molecules can have single, double or triple bonds

56. 2. Naming Molecular Compounds 1. The first element in the formula is always named first, using the entire element name. 2. The second element in the formula is named using the root of the element and adding the –ide 3. Prefixes are used to indicate the number of atoms of each type that are present (Exception: never use a prefix if there is only one of the first element)

57. NAMING COVALENT COMPOUNDS Prefixes SubscriptPrefix 1mono- 2di- 3tri- 4tetra- 5penta- SubscriptPrefix 6hexa- 7hepta- 8octa- 9nona- 10deca- Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

58. 3. WRITING MOLECULAR FORMULAS A) Write the symbols for the elements B) Place the number of the prefix after the element symbol C) Do not cross charges or reduce

59. EXAMPLES:  N 2 S 4  NI 3  XeF 6  CCl 4  P 2 O 5  SO 3 1mono 2di 3tri 4tetra 5penta 6hexa 7hepta 8octa 9nona 10deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’

60. NAMING BINARY COVALENT COMPOUNDS: EXAMPLES N 2 S 4 di nitrogen tetra sulfide NI 3 nitrogen tri iodide XeF 6 xenon hexa fluoride CCl 4 carbon tetra chloride P 2 O 5 di phosphorus pent oxide SO 3 sulfur tri oxide 1mono 2di 3tri 4tetra 5penta 6hexa 7hepta 8octa 9nona 10deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’

61. EXAMPLES:  xenon trioxide_____________________________  bromine trichloride_____________________________  nitrogen trifluoride_____________________________  tetrasulfur tetranitride_____________________________  ditellurium dichloride_____________________________  dinitrogen tetraoxide_____________________________  phosphorus pentachloride____________________________ 1mono 2di 3tri 4tetra 5penta 6hexa 7heptaa 8octa 9nona 10deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’

62.  xenon trioxideXeO 3  bromine trichlorideBrCl 3  nitrogen trifluorideNF 3  tetrasulfur tetranitrideS 4 N 4  ditellurium dichlorideTe 2 Cl 2  dinitrogen tetraoxideN 2 O 4  phosphorus pentachloride PCl 5

63. 3. Properties of Molecular compounds a. molecular compound- compound of molecules that are covalently bonded b. have low melting and boiling points c. most are gases and liquids d. do not conduct electricity