1 Chemical Bonding Chapters 8-9 (Ionic, Covalent) Chemistry
2 Answers to Explain (Analysis), Part 1: 1. Examine the data collected for melting point. What conclusions can you draw about the melting point of these chemicals?- The chemicals that took longer to melt have a higher melting point than those that melted more quickly.2. Which substances have higher melting points? Which have lower melting points? What does this indicate about the bonds in the substances? -Substances with ionic bonds had higher melting points and those with covalent bonds had lower melting points. A lower melting point indicates weaker bonds that will be more easily broken. A higher melting point indicates stronger bonds.3. Summarize the solubility of the substances in the Explore Activity.-All substances are soluble in water, except benzoic acid.
3 4. How is solubility associated with the type of bond present? -All our ionic substances are soluble in water. However, not all ionic compounds are soluble in water. One of our covalent molecules are soluble (dextrose) and one is not (benzoic acid). 5. What does the solubility of the different substances indicate about the type of bond present? -Substances with weaker intermolecular forces are more soluble in water. Those with stronger intermolecular forces are less soluble in water.6. How is conductivity related to the type of bond present?-Ionic substances, when dissolved in water, conduct electricity. Covalent substances do not conduct electricity when dissolved in water.
4 7. Why do substances with certain types of bonds conduct electricity well, while some substances are not good conductors?-Ionic compounds conduct electricity well because they possess ions, which allows electrons to flow from atom to atom. This occurs only when they are melted or dissolved in water.8. Is there a significant difference in appearance between the substances with covalent bonds and those with ionic bonds? What properties did you notice you could not see with the naked eye?-Both ionic and covalent compounds appear to be white solids. Under the hand lens, however, the covalent substances are smaller in particle size than the particles in ionic compounds. Ionic particles also have a more geometric, crystalline shape while covalent particles vary in shape.
5 -dots represent valence electrons 9. Imagine looking at the substances under a microscope. What do you think the substances might look like on a microscopic level?-(Answers will vary.)Lewis Structures are used to show bonding in molecules and ionic compounds.-dots represent valence electrons-in ionic bonding, the charge of each ion must be shown-in covalent bonding, bonded electrons is shown by lines
6 Lewis Structures: Ionic Bonds Arrows represent transfer of electrons from the metal to the nonmetal. -the charge of each atom must be shown Example: CaS
7 Lewis Structures: Single Covalent Bonds When we show bonding, shared electron pairs can be shown by either a pair of dots or a single line. -Lewis Structures are used to show how bonding electrons are arranged in molecules -example: NH3 -sigma bond (s): single covalent bond formed when an electron pair is shared by the direct overlap of orbitals ♦can occur between s & s, s & p , or p & p orbitals
8 Explain(Analysis), Part 2: Draw Lewis structures for the following ionic compounds.1. NaCl2. MgO3. LiFDraw Lewis structures for the following covalent compounds.1. H2O2. CO23. NH4
9 Lewis Structure: Multiple Bonds A multiple bond forms when two atoms share more than 2 electrons. -double bond: 4 electrons shared ( 2 pairs) ♦ O2 -triple bond: 6 electrons shared (3 pairs) ♦ N2 Some molecules have both single and multiple bonds. ♦HCN pi bond (p): forms when parallel orbitals overlap to share electrons -only occurs with multiple bonds because the first overlap is always a sigma bond
10 Lewis Structures Practice 1 Show the formation of the ionic compound for the following pairs of elements 1. strontium and fluorine 2. aluminum and oxygen 3. cesium and phosphorus 4. lead and chlorine 5. potassium and iodine 6. magnesium and chloride 7. aluminum and bromide 8. cesium and nitride 9. barium and sulfide
11 Lewis Structure Practice 2 1. PH3 2. H2S 3. HCl 4. SCl2 5. SiH4 6. CO2 7. CH2O 8. C2H2
12 Forming Chemical Bonds chemical bond: force that holds two atoms together -creates stability in the atom Two types of bonds: 1. Attraction between a positive nucleus and negative electrons (covalent bonding) 2. Attraction between a positive ion and a negative ion (ionic bonding) Remember: It is the valence electrons that are involved in this bonding.
13 Formation of Ionic Bonds ionic bond: electrostatic force that holds oppositely charged particles together -called ionic compounds -forms between metals and nonmetals ◊metals lose electrons, forms a cation ~cation: positive ion from loss of electrons ◊nonmetals gain electrons, forms an anion ~anion: negative ion formed from gain of electrons -most are binary, which means they contain 2 different elements, such as MgO, Al2O3
14 Properties of Ionic Compounds It is the chemical bonds between atoms that determines many of the physical properties of the compound.-alternating positive and negative ions form an ionic crystal-the ratio of positive to negative ions is determinedby the number of electrons transferred-strong attractionresults in a crystallattice, a 3-Darrangement ofatoms.
15 Other characteristics include: -high melting and boiling points -very hard and rigid -brittle -electrolyte when dissolved in water (aqueous solution) During chemical reactions, energy is either absorbed (endothermic) or released (exothermic) -the formation of ionic bonds is always exothermic
16 lattice energy: energy required to separate one mole of ions of an ionic compound -the more negative the lattice energy, the stronger the bondLattice Energies of Some Ionic CompoundsCompoundLattice EnergyName(kJ/mol)KI-632KF-808KBr-671AgCl-910RbF-774NaFNaI-682LiF-1030NaBr-732SrCl2-2142NaCl-769MgO-3795
17 Lattice Energyies of Some Ionic Compounds Name(kJ/mol)KI-632KF-808KBr-671AgCl-910RbF-774NaFNaI-682LiF-1030NaBr-732SrCl2-2142NaCl-769MgO-3795Depends on: 1. smaller ions -more negative value because the attraction is stronger between the nucleus and valence electrons 2. larger the positive/negative charge, the more negative the lattice energy because the attraction is stronger when more electrons are lost/gained
18 Ionic Bonding Review 1 (finish for HW) 1. Draw the Lewis dot notation showing the bonding between beryllium and chlorine. 2. What determines the properties of an element? 3. What is a crystal lattice? 4. List 5 characteristics of ionic compounds. 5. What is the difference between endothermic and exothermic? Which occurs in ionic reactions? 6. What is lattice energy? 7. What does lattice energy depend on? 8. Which substance has a stronger bond: NaCl or NaBr? Why?
19 Covalent Bonds (9.1)Remember that atoms bond to increase stability, which occurs when an atom gets a full outer shell of electrons. -in ionic bonding, one atom loses electrons (metal) and another gains electrons (nonmetal) to form oppositely charged ions with a full outer shell However, sometimes there is not a transfer of electrons, but a sharing of electrons. -covalent bond: attractive force between atoms due to the sharing of valence electrons
20 Covalent bonds can form between: -2 or more nonmetal atoms-metalloids (especially the ones to the right of themetalloid line) and nonmetalsmolecule: when two or more atoms bond covalentlyCovalent bonds can have either single bonds or multiple bonds.-single bonds: 2 shared electrons (1 pair)-multiple bonds: 4 or 6 electrons shared (2 pair=double or 3 pair = triple)
21 Properties of Molecules (Covalent Compounds) 1. low melting and boiling points. 2. many vaporize readily at room temperature 3. relatively soft solids (but not all, some are gases/liq.) 4. can form weak crystal lattices 5. do not conduct electricity when dissolved in water
22 Properties of Molecules These properties are due as a result of differences in attractive forces-attraction between atoms within a molecules is strong-attraction between different molecules is weak~called intermolecular forces or van der Walls forcesTypes of Intermolecular Forces (van der Walls forces)dispersion force (induced dipole)dipole-dipole forcehydrogen bonding
23 Properties of Molecules dispersion force (induced dipole) -occurs between nonpolar molecules -very weak dipole-dipole force -occurs between polar molecules -the more polar the molecule, the stronger the force hydrogen bonding -strong intermolecular force between the hydrogen end of one dipole and a fluorine, oxygen or nitrogen atom on another molecule’s dipole
24 Strength of Covalent Bonds All bonds can be broken, though some more easily than others. -due to the strength of the bond What affects bond strength? bond length: distance that separates the bonded nuclei -determined by the size of the atoms and how many electron pairs are shared ♦larger the atom, the longer the bond length, the weaker the bond ♦more shared electrons gives a shorter, stronger bond
25 When a bond forms or breaks, an energy change occurs When a bond forms or breaks, an energy change occurs. -bond formation: energy released (exergonic) -bond breaking: energy absorbed (endergonic) bond dissociation energy: amount of energy required to break a specific covalent bond -always a positive number -indicates the strength of a covalent bond larger the bond dissociation energy, stronger the bond (see p 246 for examples)
26 Covalent Bonding Review Describe a covalent bond.What types of atoms do covalent bonds form between?Describe single, double and triple bonds.What do we mean by sigma and pi bonds?What do we call covalent compounds?What affects bond strength?Describe the two things that determine bond length.What does bond dissociation energy indicate?What occurs when a bond forms or breaks?
27 Electronegativity and Polarity Remember that atoms have different attractions for electrons (electronegativity). -electronegativity increases left to right and decreases down a period The character and type of bond can be predicted using the difference in electronegativities between bonded atoms. -pure covalent bond: electronegativity difference = 0 (usually occurs between identical atoms, H2)
28 Most atoms do not have equal sharing of electrons, producing a purely covalent bond. -polar covalent bond: unequal sharing of electrons ♦the larger the electronegativity difference, the more ionic the bond character -ionic bonds form when the electronegativity difference is > 1.7 and nonpolar covalent bonds form when the difference is < 0.5 -the cutoff between polar covalent, nonpolar, and ionic is sometimes inconsistent with experimental data
29 Electronegativity Practice Remember: bonding is not clearly ionic or covalent, with ionic character increasing as the difference in electronegativity increases.Decide if the following pairs of atoms are polar covalent, nonpolar covalent or ionic.N-H= 0.84polar covalentC-Cl= 0.61S-Se= 0.03nonpolar covalent
30 When a polar bond forms the shared electrons are pulled more strongly toward one atom. -this creates partial charges at opposite ends of the molecule, which is called a dipole ♦ d- indicates a partial negative d+ indicates a partial positive Polar molecule or not? A molecule can have individual polar bonds, but make a nonpolar molecule. How? We look at the shape of the molecule.
31 Let’s look at H2O and CCl4. O—H C—Cl d- d+ d+ d- 1. 24 0 Let’s look at H2O and CCl4. O—H C—Cl d- d+ d+ d both O-H and C-Cl have polar covalent bonds One molecule is polar and the other is nonpolar? How do we know? We look at the shape of the molecule and the terminal atoms.
32 -symmetric molecules like CCl4 are nonpolar because the polar bonds cancel each other out. CCl4 -asymmetric molecules like H2O are polar because the polar bonds do not cancel each other out. H2O
33 If water is polar, why will oil not dissolve in it If water is polar, why will oil not dissolve in it? Oil must be nonpolar because A substance is only soluble (dissolvable) when combined with a like molecule. “Like Dissolves Like” hydrophobic- “fear of water” hydrophilic- “likes water”
34 Polarity Review1. What is electronegativity and what does it predict? 2. What is the difference between a nonpolar covalent bond and a polar covalent bond? 3. What is a dipole and what indicates them? 4. Describe the electronegativity trend both across a period and down a group. 5. Are the following bonds polar or nonpolar covalent? a. H-Br b. C-O c. S-C 6. Describe the relationship between polarity and solubility. 7. What do we mean by symmetric and asymmetric?
35 Final Bonding Questions: Draw a table comparing the properties of ionic andcovalent bonds.-leave room to add more properties (we will discussthe table and add more to it)What is a general definition of a bond?What are the two types of bonds? Describe each.What is the octet rule?What do we mean by polar or nonpolar?What is electronegativity? How do we use this in bonding?What are intermolecular forces?
36 1. 2. A bond is a force holding two atoms together to create stability in an atom 3. An ionic bond is an attraction of oppositely charged ions due to a transfer of electrons from a metal atom to a nonmetal atom. A covalent bond is the sharing of electrons between nonmetals or nonmetals and some metalloids.IonicCovalenthigh melting/boiling pointlow melting/boiling pointelectrolyte in waternonelectrolyte in watercrystal lattice structuresome form weak crystal latticeshard, brittle solidsgases, liquids, relatively soft solidsmost dissolve in watersome dissolve in watermany vaporize at room temp
37 The octet rule states that atoms are stable if they have a full valence shell of electrons. For most atoms, the number is 8, but the period 1 elements are stable with 2.A covalent bond is polar if there is an unequal sharing of electrons due to the electronegativity difference between the atoms. It is nonpolar if there is an equal sharing of electrons.Electronegativity is the attraction an atom has for electrons. The more electronegative the atom, the stronger the attraction. We use electronegativity to determine the polarity of molecules.Intermolecular forces are the force that holds atoms together. They can be weak, allowing atoms to be pulled apart easily, or strong.
39 Molecular Structures (9.3) structural formula: uses letter symbols and bonds to show relative positions of atoms -one of the most useful -can be predicted for many molecules by drawing Lewis structures -H is always an end (terminal) atom, never a central atom -less electronegative atom is the central atom (nm or metalloid closest to the left of the PT-usually)
40 Structural Formulas-Example CH2O 1. Predict the location of the atoms C is least electronegative & farthest to left on PT, therefore it is the central atom 2. Find the total number of electrons available for bonding. 1 C-4, 2 H-2, 1 O-6 for a total of 12 valence e- 3. Determine the number of bonding pairs 12 valence e- / 2 = 6 electron pairs
41 central atom and each terminal atom. H C O H 4. Place one bonding pair (single bond) between thecentral atom and each terminal atom.H C OH5. Subtract the number of pairs you used in step 4 fromthe number of bonding pairs determined in step 3.6 – 3 used = 3 e- pairs left
42 5. Subtract the number of pairs you used in step 4 from the number of bonding pairs determined in step 3.-take the remaining electron pairs and place electronpairs around the terminal atoms to satisfy the octetruleH C OH
43 6. If the central atom is not surrounded by 4 electron pairs, it does not have an octet-convert one or two of the lone pairs on a terminalatom to a double or triple bond between that terminalatom and the central atomH C OHPractice:1. CH3Cl NBr5
44 Structural Formulas-Polyatomic Ions Writing structural formulas for polyatomic ions is the same with one exception: -the total number of electrons may differ due to the negative and positive charge. ♦negative charge, more electrons are present SO4-2 add two electrons ♦positive charge, less electrons are present NH4+1 subtract one electron
45 Resonance StructuresLet’s look at CO3-2. -when one or more valid Lewis structure can be written for a molecule, resonance occurs -let’s look at another resonance molecule/ion: NO3-1 -each molecule/ion that undergoes resonance behaves as if it only has one Lewis structure
46 Exceptions to the Octet Rule Some molecules do not obey the octet rule. Three reasons exist: 1. when a small group of molecules have an odd number of valence electrons: -NO2 for a total of 17 valance electrons-one unpaired electron on N
47 2. Some form with fewer than eight, though this is relatively rare: -B in BH3 is stable with six because it only has 3 valence electrons. 3. When the central atom has more than 8 electrons, which is referred to as an expanded octet. -can occur in elements that are found in period three or higher elements (because of the d orbitals). -P in PCl5 (1 s orbital, 3 p orbitals, and 1 d orbital)
48 Structural Formulas Practice 1. SO3 2. N2O 3. SF6 4. ClF3 5. SiF4 6. PO BF3 8. SO3-2
49 Molecular Structure Review 1. What is a structural formula? 2. Describe resonance. 3. List three reasons for exceptions to the octet rule. 4. Name the following: a. BH3 b. SO2 c. PO Write formulas for the following: a. sulfur trioxide c. chlorous acid b. hydrosulfuric acid 6. Draw structural formulas a. SO2 b. H2O c. BrCl5
50 Molecular ShapeMany of the physical and chemical properties of molecules is determined by the shape of the molecule. -the shape of molecules determines if two or more molecules can get close enough for a reaction to occur. VSEPR (Valence Shell Electron Pair Repulsion) model: atoms in a molecule are arranged so that the pairs of electrons (bonded and lone) minimize repulsion.
51 VSEPR modelThe repulsion between electron pairs result in fixed angles between atoms -bond angle: angle formed by any two terminal atoms and the central atom ♦lone pairs take up slightly more space than bonded pairs ♦multiple bonds have no affect on the geometry because they exist in the same region as single bonds -example: H2O See page 260 for the Molecular Geometries (Shapes)
52 Shapes Review 11. What determines many of the physical and chemical properties of molecules? 2. Describe the VSEPR model. 3. What does the repulsion between electron pairs result in? 4. Why do multiple bonds have no affect on geometry of a molecule? 5. Why do molecules with lone pairs have shorter bond angles? 6. How do we know if a molecule is polar or nonpolar?
53 Names and Formulas-Ionic Compounds A universal set of rules must be used so chemists around the world can communicate. formula unit: simplest ratio of ions represented in an ionic compound -remember that ionic compounds form a crystal lattice, consisting of many cations and anions. -the overall charge for the compound is 0 Most ionic compounds are binary, consisting of two monatomic ions. -monatomic ion: one atom ion, either positively or negatively charged
54 Remember that we determine the charge of each ion by its oxidation number. Formula Rules for Ionic Compounds 1. write the cation first, followed by the anion 2. state the charges of both ions 3. cross the number for the charge of one ion to become the subscript for the other ion. -subscripts are used to state the number of each atom in the compound
55 Example: Determine the formula for the ionic compound formed when potassium reacts with oxygen. 1. Cation = potassium = K Anion = oxygen = O 2. K+1 O-2 3. K+1 O-2 K2O1 K2O You try: Determine the formula for the ionic compound formed when aluminum reacts with chlorine.
56 Ionic Compounds with Polyatomic Ions We write formulas for ionic compounds containing polyatomic ions the same way as in binary compounds. -the cation comes first, followed by the anion -state the charges -cross over the number for the charges However: -if you have more than one polyatomic ion, place parenthesis around the polyatomic ion, with the subscript outside the parenthesis.
57 Example: Determine the formula for the ionic compound formed when beryllium reacts with cyanide. 1. Cation = beryllium = Be Anion = cyanide = CN- 2. Be+2 CN-1 3. Be+2 CN-1 Be1(CN)2 Be(CN)2 You try: Determine the formula for the ionic compound formed when ammonium reacts with iodine.
58 Ionic Bonding Practice 2 Write the correct formula for the following pairs of atoms: 1. ammonium and oxygen 2. lithium and nitrate 3. aluminum and hydroxide 4. ammonium and phosphate 5. strontium and acetate
59 Ionic Bonding Review 21. Why do we need a universal set of rules for naming and writing formulas? 2. Define monatomic and binary. 3. What is meant by a formula unit? 4. What is the purpose of subscripts. 5. Describe what a polyatomic ion is? 6. When do we use parenthesis for writing ionic compounds with polyatomic ions? 7. Determine the formula for the ionic compound formed when lead reacts with sulfur. 8. Determine the formula for the ionic compound formed when lithium reacts with nitrogen.
60 Ionic Bonding Practice 3 Write the correct formula for the following pairs of atoms: 1. aluminum and carbon 2. ammonium and carbonate 3. calcium and oxygen 4. aluminum and chromate 5. sodium and phosphate 6. potassium and hydrogen sulfate 7. magnesium and phosphorus
61 Naming Ionic Compounds The names of ionic compounds include the ions of which they are composed. 1. The element whose symbol appears first in the formula also appears first in the name. -this is always the positively charged ion, or metal 2. The name of the second ion follows, with its ending changed to –ide for single atom ions. Ex: What is the name of MgCl2? magnesium chloride
62 Ionic Compounds Practice 4 Write the formula and the name.1. Na2S2. Ga2S33. CaSe4. LiF
63 Naming with Polyatomic Ions You follow the same rules when naming polyatomic ions as when you have binary ionic compounds, however: -you do not change the ending of the polyatomic ions, even when they are the second atom. Example: Al2(SO4)3 aluminum (III) sulfate Rule: You must state the charge of all metals not included in groups 1 and 2 because many have multiple charges.
64 Ionic Compounds Practice 5 Name the following compounds: 1. NaC2H3O2 2. CaCO3 3. KOH 4. Mg(NO3)2 5. Li2CrO4 6. Mg(NO2)2 7. AlPO4
65 Rules for Transition Metals *According to the previous rules, FeO and Fe2O3 would both be named iron oxide,even though they are not the same compound* Since many transition metals can have more than one charge, the name must show this. This is done using roman numerals. -FeO is named iron (II) oxide because Fe has a +2 charge -Fe2O3 is named iron (III) oxide because Fe has a +3 charge *The roman numeral states the charge of the metal*
66 Q: How do I know the iron in FeO has a +2 charge Q: How do I know the iron in FeO has a +2 charge? A: The oxide ion has a –2 charge, so the Fe must have a +2 charge so the compound is overall neutral. Q: How do I know the iron in Fe2O3 has a +3 charge? A: There are three oxide ions with a –2 charge: (3 ions)(-2 charge/ion) = a total of –6 charge Since the overall charge must be neutral, the iron must have a total charge of +6. Therefore: (2 ions)(x charge/ion) = +6 x = +3
67 Ionic Compounds Practice 6 Write the formula given & the name of each compound. 1. FeCl3 2. Zn3P2 3. CuS 4. AuF 5. CuC2H3O2 6. AgHCO3 7. ZnSO4 8. Pb(CO3)2
68 Ionic Bonding Review 31. What is the ending of the second atom changed to when naming ionic compounds? 2. Write the name for (NH4)3P 3. Write the name for AlS. 4. Determine the formula for the ionic compound formed when magnesium reacts with phosphate. 5. Determine the formula for the ionic compound formed when magnesium reacts with phosphate.
69 Naming Molecules (9.2)Molecules are represented by both names and formulas. Rules for Naming Binary Molecular Compounds 1. The first element in the formula is 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 suffix –ide. 3. Prefixes are used to indicate the number of atoms of each type that are present in the compound. -exception: 1st element never uses the prefix mono- -drop the final letter of the prefix if element name begins with a vowel.
70 Prefixes you need to know: # atoms prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-
71 Naming Binary Molecules-Example Name the compound P2O5, which is used as a drying and dehydrating agent. 1st atom: P = phosphorus 2nd atom: O = oxygen = oxide There are 2 phosphorus = diphosphorus There are 5 oxygens = pentoxide (drop the –a of penta-) Put it together: diphosphorus pentoxide
72 Naming Binary Molecules Practice Name the following molecules: 1. CCl4 2. As2O3 3. CO 4. SO2 5. NF3
73 Naming Acids(We will talk more about acids in Ch 19) There are two types of acids: 1. binary acid: contains hydrogen and one other element -when naming use the prefix hydro- plus the root of the second element with the suffix –ic, followed by the word acid. -ex: HCl H = hydro- Cl = chloride = chloric hydrochloric acid
74 Some acids are not binary, but are named according to the binary acid rules when oxygen is not present, as in HCN. H = hydro CN = cyanide = cyanic hydrocyanic acid 2. oxyacid: an acid that contains an oxyanion (oxygen containing polyatomic ion) -the name depends on the oxyanion present -the name consists of the root of the anion, a suffix, and the word acid ♦if the anion suffix is –ate, it is replaced with -ic ♦if the anion suffix is –ite, it is replaced with -ous
76 Naming Acids PracticeName the following acids: 1. HBr 2. H3PO4 3. H2SO4 4. H2SO3 5. H2CO3
77 Writing FormulasUse the prefixes in the molecule’s name to determine the subscript for each atom in the compound. - phosphorus tribromide P Br 1 (no prefix) 3 (tri) PBr3 - the formula for an acid can be derived from the name as well ♦charge of the oxyanion or anion gives the number of hydrogens hydrofluoric acid = HF (1 H because fluorine has a -1 charge)
80 Metallic BondsMetallic bonds are similar to ionic bonds because they often form lattices in the solid state. -eight to twelve metal atoms surround another, central metal atom Instead of sharing electrons or losing electrons, the outer orbitals overlap. -electron sea model: all metal atoms in a metallic solid contribute their valence electrons to form a ‘sea’ of electrons around the metal atoms. -valence electrons are free to move from atom to atom (delocalized electrons), forming metallic cations
82 metallic bond: attraction of a metallic cation for the delocalized electrons that surround it This bonding contributes to the unique properties of metals: 1. generally have high melting and boiling points, with especially high boiling points -due to the amount of energy needed to separate the electrons from the group of cations 2. malleable (hammered into sheets) and 3. ductile (drawn into wire) -mobile electrons can easily be pulled and pushed past each other
83 4. durable-though electrons move freely, they are stronglyattracted to the metal cations and are not easilyremoved from the metal
84 5. good conductors -free movement of the delocalized electrons, allowing heat and electricity to move from one place to another very quickly 6. luster -interaction between light and delocalized electrons
85 As the number of delocalized electrons increases, as in transition metals (d electrons), the hardness and strength also increases. -alkali and alkaline earth metals are soft (s valence electrons only) It is easy to combine 2 or more different metals to make a metallic crystal -alloy: mixture of elements with metallic properties -the properties of alloys differ from those of the individual elements that make it up
86 Metallic Bonding Bellringer What is a metallic bond?What is an alloy?Describe the electron sea model.What occurs with orbitals in metals?How is metallic bonding similar to ionic bonding?What are delocalized electrons?What contributes to a metal’s high boiling point, malleability, ductility and conductivity?List the other 2 properties of metals.What happens to strength and hardness as you decrease the number of delocalized electrons?