Ionic, Covalent, Metallic

Transcription

1 Ionic, Covalent, Metallic

2 Physical Properties of Types of Compounds IONIC COVALENT METALLIC Attractive/force strength Melting/Boiling point Strong Weak Varies High Low Varies Vapor pressure Low High Varies Electrical conductivity Only in liquid or aquous solution Never! Always! Malleable/Brittle Brittle Brittle Malleable Type of bond M + NM NM + NM M Example NaCl,. ZnSO 4 H 2 O, NH 3, C 6 H 12 O 6 Zn, Cu, Ca, Mg, Na, Fe

3 Naming Ionic Compounds

4 Naming Ionic Compound Metal is listed first, followed by nonmetal. Change the name of the nonmetal to -ide. Examples: nitride, sulfide, fluoride, oxide, bromide, iodide, chloride, telluride, phosphide.

5 The 5 Steps for writing an ionic compound formula 1. Write the symbols of the two elements. 2. Write the charge of each as superscripts. 3. Drop the positive and negative signs. 4. Crisscross the superscripts so they become subscripts. (1 is not written) 5. Reduce when possible.

6 Formula for boron oxide 1. Write the symbols of the two elements. 2. Write the charge for each element. 3. Drop the positive & negative sign. 4. Crisscross the superscripts so they become subscripts. (1 is not written) 5. Reduce subscripts when possible. (not possible here) B +3 3 O

7 Formula for Aluminum sulfide 1. Write the symbols of the two elements. 2. Write the charge for each element. 3. Drop the positive & negative sign. 4. Crisscross the superscripts so they become subscripts. (1 is not written) 5. Reduce subscripts when possible. (not possible here) Al +3 3 S

8 Formula for Calcium oxide 1. Write the symbols of the two elements. 2. Write the charge for each element. 3. Drop the positive & negative sign. 4. Crisscross the superscripts so they become subscripts. (1 is not written) 5. Reduce subscripts when possible. Ca CaO O 2 2

9 Examples of Reduction of Subscripts:

10 Transition Metals Most Transition metals have two valences. Roman numerals are used in the name of the transition metal in the compound to show the valence on the cation.

11 Formula for Copper (II) Chloride 1. Write the symbols of the two elements. 2. Write the charge for each element. 3. Drop the positive & negative sign. 4. Crisscross the superscripts so they become subscripts. (1 is not written) 5. Reduce subscripts when possible. (not possible here) Cu Cl

12 The 5 Steps for writing an ionic compound formula with polyatomic ions 1. Write the symbols of the cation and anion. Look up in the polyatomic chart if needed 2. Write the charge of each as superscripts. 3. Drop the positive and negative signs. 4. Crisscross the superscripts so they become subscripts. (1 is not written) If more than one of the polyatomic ion is needed, use paranthesis to represent as a group 5. Reduce when possible.

13 Formula for Magnesium carbonate 1. Write the symbols of the cation and anion. Look up in the polyatomic chart if needed 2. Write the charge of each as superscripts. 3. Drop the positive and negative signs. 4. Crisscross the superscripts so they become subscripts. (1 is not written) If more than one of the polyatomic ion is needed, use paranthesis to represent as a group 5. Reduce when possible. MgCO ( ) 2 2

14 Formula for Zinc hydroxide 1. Write the symbols of the cation and anion. Look up in the polyatomic chart if needed 2. Write the charge of each as superscripts. 3. Drop the positive and negative signs. 4. Crisscross the superscripts so they become subscripts. (1 is not written) If more than one of the polyatomic ion is needed, use paranthesis to represent as a group 5. Reduce when possible. Zn(OH) ( ) 2 2

15 Formula for Aluminum Phosphate 1. Write the symbols of the cation and anion. Look up in the polyatomic chart if needed 2. Write the charge of each as superscripts. 3. Drop the positive and negative signs. 4. Crisscross the superscripts so they become subscripts. (1 is not written) If more than one of the polyatomic ion is needed, use paranthesis to represent as a group 5. Reduce when possible AlPO 3 4 ( ) 3 3

16 Steps to Ionic Formula to Naming 1. Identify the cation: The cation is always written first in the formula for an ionic compound. The only common polyatomic ion that you will encounter is the ammonium ion (NH 4+ ). 2. Identify the anion: Cover up the cation. Everything that is leftover will be the anion. 3. Write the name of the ionic compound by writing the name of the cation followed by the name of the anion. If polyatomic, use the chart to find its charge If transition metal, calculate the charge by balancing it out with anion s charge

17 Write the correct name for (NH 4 ) 2 CO Identify the cation: The cation will always be the first ion written. 2. Identify the anion: The anion will be everything leftover once the cation has been identified: (NH 4 ) 2 CO 3 carbonate ion 3. The correct name for this compound is ammonium carbonate. (NH ) CO AMMONIUM CARBONATE

18 Write the name for SnO Identify the cation: The cation will always be the first ion written. 2. Identify the anion: The anion will be everything leftover once the cation has been identified 3. If transition metal, calculate The charges the charge of the by Tin balancing is a it out with anion s cationcharge and the anions transition metal, needs special charge written with +4 SnO 2 Tin (IV) Oxide Oxide Tin? the name must exactly balance out. Oxygen Oxygen

19 Write the name for Mn 3 (PO 4 ) Identify the cation: The cation will always be the first ion written. 2. Identify the anion: The anion will be everything The leftover charges once of the the cation has been identified cation and the anions must exactly balance 3. If transition metal, calculate the charge by Tin balancing is a it out with anion s out. charge transition Mn (PO )-6 metal, Manganese needs special charge? -3 written +2 with Manganese the name 3? Manganese +2 Phosphate (PO ) 4 2 Phosphate Manganese Phosphate? Manganese (II) Phosphate

20 Write the name for Cu (OH) Identify the cation: The cation will always be the first ion written. 2. Identify the anion: The charges of the cation and the anions must exactly balance The anion will be everything leftover once the cation has been identified: 3. If transition metal, calculate the charge by out. Copper balancing is a it out with anion s charge +3-3 transition metal, needs Cu special charge (OH) Cupper written with? the name 3-1 Copper Hydroxide Copper (III) Hydroxide Hydroxide -1 Hydroxide -1 Hydroxide

21 Checking for understanding Name these ionic compounds: KI CaBr₂ Na₃N Ca(OH)₂ Write formulas: Barium oxide Calcium phosphate Sodium hydroxide Magnesium sulfide

22 Click Below for the Video Lectures Ionic Solids Ionic Bonding

23 Naming Covalent Molecules

24 Covalent Molecule Naming Rules Compounds between two nonmetals 1. First element in the formula is named first. Keeps its element name Gets a prefix if there is a subscript on it 2. Second element is named second Use the root of the element name plus the -ide suffix Always use a prefix on the second element

25 List of Prefixes 1 = mon(o) 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deka

26 Naming P 2 O 5 1. First element in the formula is named first. Keeps its element name Gets a prefix if there is a subscript on it 2. Second element is named second Use the root of the element name plus the -ide suffix Always use a prefix on the second element P O 2 5 Diphosphorus O 5 Phosphorus pentoxide Oxide

27 Naming Covalent Compounds P 2 O 5 = CO 2 = CO = N 2 O = diphosphorus pentoxide carbon dioxide carbon monoxide dinitrogen monoxide 10/25/

28 Checking for understanding Name these covalent compounds: CO PBr₅ H₂O SeF₆ Write formulas: Carbon Tetrachloride Ammonium Sulfate Dinitrogen Trifluoride Pentaphosphorus Heptoxide

29 Click Below for the Video Lectures Covalent Bonding Covalent Network Solids Molecular Solids

30 Drawing Covalent Molecules

31 The Octet Rule: The Diatomic Fluorine Molecule F F 1s 2s 2p Each has seven valence electrons 1s 2s 2p F F

32 The Octet Rule: The Diatomic Oxygen Molecule O O 1s 2s 2p Each has six valence electrons 1s 2s 2p O O

33 The Octet Rule: The Diatomic Nitrogen Molecule N N 1s 2s 2p Each has five valence electrons 1s 2s 2p N N

34 Drawing Covalent Molecules 3. CH 3 NH 2 (14 e- ) 1. HBr 2. NH 3 (8 e- ) (8 e- ) H H H Br H N H H C N H H H Needs 2 e- 1. Count the total number of valence electrons 2. Place the atom that makes most bonds in the middle. (least electronegative other than hydrogen) 3. Draw single bonds to the other atoms off of the central atom. 4. Place electrons around peripheral atoms to fill octet, then to central atom H Has 2 e- H H C N H H 5. if electrons run out, start making double, triple bonds 34

35 1. Count the total number of valence electrons 2. Place the atom that makes most bonds in the middle. (least electronegative other tha hydrogen) 3. Draw single bonds to the other atoms off of the central atom. 4. Place electrons around peripheral atoms to fill octet, then to central atom 5. if electrons run out, start making double, triple bonds 4. CH4 H H C H (8 e- ) 5. O3 (18 e- ) H O Use these two to make a double bond O O 2 electrons deficient 6. CO2 (16 e- ) O C O 4 electrons deficient O C O O O O 10/25/

36 Molecular Force

37 Nonpolar molecules form LONDON DISPERSION FORCE attractions. Since there are no permanent positive or negative ends, these attractions are extremely weak. The attractions are a combination of temporary poles due to electron movement around the molecule or, in the case of huge molecules, they actually get tangled up with each other like sticky strands of spaghetti or yarn. London dispersion forces generally get stronger as the size of the molecule increases.

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39 Polar molecules form DIPOLE attraction: the simple attraction of the oppositely charged ends of two molecules. The partially positive end of one molecule attracts to the partially negative end of a different molecule. This attraction allows these substances to exist as solids and liquids at higher temperatures than are possible for nonpolar molecules of equivalent size.

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41 Hydrogen Bond Some polar molecules form HYDROGEN BONDS between them. Bonding between hydrogen and more electronegative neighboring atoms; Nitrogen, Oxygen and Flourine (F,O,N) An extremely high melting point, boiling point, heat of fusion and heat of vaporization for a molecule its size

42 Hydrogen Bonding in Water The attraction between the two molecule is nearly that of ionic attraction. Not only that but the H end of one molecule actually forms temperary covalent bonds with the N,O or F that makes up the end of the other molecule. This gives extra strength that allows water to be liquid at room teperature despite its small size.

43 Hydrogen bonding in water

44 So, what should you be able to do now? 1) Identify whether a compound is molecular, ionic, metallic or network based on its properties 2) Draw dot diagrams of simple molecules 3) Draw structural formulas of simple molecules 4) Determine the shape of simple molecules 5) Determine if simple molecules are polar or nonpolar 6) 7) Determine the attractive force type that attracts specific simple molecules to each other.

45 Checking for understanding 1. Explain the difference between polar and nonpolar molecules 2. Explain London dispersion force 3. Explain dipole attraction forces including hydrogen bond

46 Click Below for the Video Lectures Intermolecular forces Dipole Forces London Disperson forces

47 Chemical Formula

48 Click Below for the Video Lectures Molecule and Elements

49 Molecular Formula Molecular Formula - indicates the total number of atoms of each element needed to form the molecule. MOLECULES ARE PARTICLES FORMED FROM THE COVALENT BONDING OF NONMETAL ATOMS. A molecule of methane contains one atom of carbon and four atoms of hydrogen, so the molecular formula is CH4.

50 Empirical Formula Empirical Formulas represents the simplest ratio in which the atoms combine to form a compound. Empirical formulas are used to represent ionic compounds which form crystals of alternating + and charge instead of separate molecules. CaCl 2 is ionic, and the formula represents a ratio of 1 ion of Ca +2 to every 2 ions of Cl -

51 Interpreting chemical formula The subscript tells you how many moles of that particular ion you have in one mole of the compound. A Mole is to chemistry what a dozen is to donuts or eggs. Chemical formulas and reactions are written by the mole. 1 mole = 6.02 X of anything. If there is no subscript, then the number of moles of that ion is 1. When Interpreting Formulas, if there are (parentheses) around the element, any subscript outside the parentheses multiplies all of the elements inside the parentheses by that amount.

52 Ca(NO ) outside the parentheses, so double the number of atoms inside to get the total number of atoms of that element (2 N s and 6 O s). Since Ca is not inside the parentheses, it is not doubled. There is only 1 Ca in the formula. 1 mol Ca, 2 mol N, 6 mol O 6.02x10 23 atoms of Ca 2 x (6.02x10 23 ) atoms of N 6 x (6.02x10 23 ) atoms of O

53 Checking for understanding Complete the chart bellow, an example was done for you BaCO 3 1mol Ba 1mol C 3mol O Na 3 PO 4 Al 2 (SO 4 ) 3

54 Molar Mass

55 Molar Mass The molar mass is the mass of one mole of an element or compound. Expressed in g/mol 55

56 Molar Mass from Periodic Table 1 mole Ag 1 mole C 1 mole S = g = g = g

57 Molar Mass of an Element To calculate molar mass use the atomic masses of all elements present in a compound Molar Mass of CO 2 = 1 x g/mol + 2 x g/mol = g/mol

58 Molar Mass of a Compound The molar mass of a compound is the sum of the molar masses of the elements in the formula. Example: Calculate the molar mass of CaCl 2. Element Number of Moles Atomic Mass Total Mass Ca g/mole 40.1 g Cl g/mole 71.0 g CaCl g 58

59 Molar Mass of K 3 PO 4 Calculate the molar mass of K 3 PO 4. Element Number of Moles Atomic Mass Total Mass in K 3 PO 4 K g/mole g P g/mole 31.0 g O g/mole 64.0 g K 3 PO g 59

60 Mole Conversion

61 Mole Map 6.02x10 23 Molar Mass atoms moles grams 1. Circle the numbers 2. Underline what you are looking for 3. Start with what you circled

62 Mass to Mole Conversion How many moles of carbon is 26 g of carbon? 26 g C 1 mol C g C = 2.2 mol C moles Molar Mass grams 62

63 How many moles is 5.69 g of Na? 5.69 g Na 1 mol Na g Na = mol Na moles Molar Mass grams 63

64 How many grams are in 9.45 mol of nitrogen atoms? 9.45 mol N g N 1 mol N = =132 g N moles Molar Mass grams 64

65 Remember 1 Mole = 6.02 x atoms or molecules 0.5 Mole = 3.01 x atoms or molecules 2 Mole = 1.20 x atoms or molecules

66 Molar Conversion Examples How many atoms are in 2.50 moles of C? 2.50 mol atoms 1 mol =atoms C atoms 6.02x10 23 moles 66

67 Calculate the number of atoms in mol of Al. 0.5 mol atoms 1 mol =atoms Al atoms 6.02x10 23 moles 67

68 Calculate the number of moles of 1.80 x Na atoms x atoms 1 mol atoms = 2.99 moles atoms 6.02x10 23 moles

69 How many grams are in 1.20 x10 24 molecules of dinitrogen trioxide, N 2 O 3? molecules 1 mol molecules g 1 mol = g =152 g N 2 O 3 atoms 6.02x10 23 moles Molar Mass 69 grams 69 69

70 Find the mass of molecules of NaHCO molecules 1 mol molecules g 1 mol = 290 g NaHCO 3 atoms 6.02x10 23 moles Molar Mass 70 grams

71 Checking for understanding 1. How many atoms are in 5g of Ca? 2. How many atoms are in 1.2 mol K? 3. Calculate the mass of 1.4x10 24 molecules of CaCO 3

72 Click Below for the Video Lectures The Mole