Inorganic Chemistry Chapter 11

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Gas
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Expand to fill their container
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Liquids
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Retain volume, but not shape
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Solids
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Retain volume and shape
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Physical state of molecule depends on..
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▪ Average kinetic energy of particles ▪ Intermolecular Forces
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Physical properties of gases, liquids and solids determined by..
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▪ How tightly molecules are packed together ▪ Strength of attractions between molecules
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Converting gas → liquid or solid
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▪ Molecules must get closer together ▪ Cool or compress
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Converting liquid or solid → gas
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▪ Requires molecules to move farther apart ▪ Heat or reduce pressure
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Intramolecular forces
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▪ Covalent bonds within molecule ▪ Strong
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Intermolecular forces
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▪ Attraction forces between molecules ▪ Weak
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Electronegativity
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Measure of attractive force that one atom in a covalent bond has for electrons of the bond
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Bond Dipoles
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▪ Two atoms with different electronegativity values share electrons unequally ▪ Electron density is uneven ▪ Higher charge concentration around more electronegative atom ▪ Indicated with delta (δ) notation
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Symmetrical molecules
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▪ Even if they have polar bonds ▪ Are non-polar because bond dipoles cancel
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When substance melts or boils..
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▪ Intermolecular forces are broken ▪ Not covalent bonds
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Asymmetrical molecules
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▪ Are polar because bond dipoles do not cancel ▪ These molecules have permanent, net dipoles
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Molecular dipoles
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▪ Cause molecules to interact ▪ Decreased distance between molecules increases amount of interaction
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What are intermolecular forces responsible for?
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▪ Responsible for non-ideal behavior of gases ▪ Responsible for existence of condensed states of matter ▪ Responsible for bulk properties of matter ▪ Boiling points and melting points ▪ Reflect strength of intermolecular forces
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Types of intermolecular forces?
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1. Dipole-dipole forces ▪ Hydrogen bonds 2. London dispersion forces 3. Ion-dipole forces ▪ Ion-induced dipole forces
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Dipole-Dipole Attractions
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▪ Occur only between polar molecules ▪ Possess dipole moments ▪ Molecules need to be close together ▪ Polar molecules tend to align their partial charges ▪ Positive to negative ▪ As dipole moment increases, intermolecular force increases ▪ Mixture of attractive and repulsive dipole-dipole forces. ▪ Decrease as molecular distance increases ▪Forces increase with increasing polarity
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Hydrogen Bonds
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▪ Special type of dipole-dipole Interaction ▪ Very strong dipole-dipole attraction ▪ Occurs between H and highly electronegative atom (O, N, or F) ▪ H—F, H—O, and H—N bonds very polar ▪ Element’s small size, means high charge density ▪ Positive end of one can get very close to negative end of another ▪ Responsible for expansion of water as it freezes ▪ Produces strong attractions in liquid
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London Forces
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▪ When atoms near one another, their valence electrons interact ▪ Repulsion causes electron clouds in each to distort and polarize ▪ Instantaneous dipoles result from this distortion ▪ Effect enhanced with increased volume of electron cloud size ▪ Effect diminished by increased distance between particles and compact arrangement of atoms ▪ Instantaneous dipole-induced dipole attractions ▪ Operate between all molecules (nonpolar or polar)
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Ease with which dipole moments can be induced and thus London Forces depend on..
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1. Polarizability of electron cloud 2. Points of attraction ▪ Number atoms ▪ Molecular shape (compact or elongated)
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Polarizability
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▪ Ease with which the electron cloud can be distorted ▪ Larger molecules often more polarizable ▪ Larger number of less tightly held electrons ▪ Magnitude of resulting partial charge is larger ▪ Larger electron cloud
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Molecular Shape
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▪ Increased surface area available for contact = increased London forces ▪ London dispersion forces between spherical molecules are lower than chain-like molecules ▪ More compact molecules ▪ Hydrogen atoms not as free to interact with hydrogen atoms on other molecules ▪ Less compact molecules ▪ Hydrogen atoms have more chance to interact with hydrogen atoms on other molecules
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Ion-Dipole Attractions
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▪ Attractions between ion and charged end of polar molecules ▪ Attractions can be quite strong as ions have full charges ▪ Negative ends of water dipoles surround cation ▪Positive ends of water dipoles surround anion ▪ Attractions between ion and dipole it induces on neighboring molecules ▪ Depends on ▪ Ion charge and ▪ Polarizability of its neighbor ▪ Attractions can be quite strong as ion charge is constant, unlike instantaneous dipoles of ordinary London forces
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Strongest to weakest intermolecular forces
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1. Ion-Dipole 2. Hydrogen Bonding 3. Dipole-Dipole 4. London Forces
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Compressibility
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▪ Measure of ability of substance to be forced into smaller volume ▪ Determined by strength of intermolecular forces
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Diffusion
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▪ Movement that spreads one gas though another gas to occupy space uniformly ▪ Spontaneous intermingling of molecules of one gas with molecules of another gas
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Intermolecular Forces and Temperature
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▪ Decrease with increasing temperature ▪ Increasing kinetic energy overcomes attractive forces ▪ If allowed to expand, increasing temperature increases distance between gas particles and decreases attractive forces
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Surface Tension
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▪ Liquids containing molecules with strong intermolecular forces have high surface tension ▪ Increases as intermolecular forces increase ▪ Decreases as temperature increases
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Wetting
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▪ Ability of liquid to spread across surface to form thin film ▪ Greater similarity in attractive forces between liquid and surface, yields greater wetting effect ▪ Occurs only if intermolecular attractive force between surface and liquid about as strong as within liquid itself
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Viscosity
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▪ Resistance to flow ▪ Measure of fluid’s resistance to flow or changing form ▪ Related to intermolecular attractive forces ▪ Also called internal friction ▪ Depends on intermolecular attractions ▪ Decreases when temperature increases
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Solid → Gas
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▪ Sublimation ▪ Endothermic
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Gas → Liquid
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▪ Condensation ▪ Exothermic
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Rate of Evaporation Depends on..
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▪ Temperature ▪ Surface area ▪ Strength of intermolecular attractions
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System at Equilibrium
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▪ Rate of evaporation = rate of condensation ▪ Occurs in closed systems where molecules cannot escape
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Endothermic sequence
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(absorbs heat) heat solid → melt → heat liquid → boil → heat gas
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Exothermic sequence
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(release heat) cool gas → condense → cool liquid → freeze → cool solid
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Endothermic Phase Changes
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1. Must add heat 2. Energy entering system (+) Sublimation: ΔHsub > 0 Vaporization: ΔHvap > 0 Melting or Fusion: ΔHfus > 0
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Exothermic Phase Changes
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1. Must give off heat 2. Energy leaving system (-) Deposition: ΔH < 0 = -ΔHsub Condensation: ΔH < 0 = -ΔHvap Freezing: ΔH < 0 = -ΔHfus
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▪Liquid-Vapor Equilibrium
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▪ Molecules in liquid ▪ Not in rigid lattice ▪ In constant motion ▪ Denser than gas, so more collisions
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Vapor Pressure
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▪ Pressure molecules exert when they evaporate or escape into gas (vapor) phase ▪ Pressure of gas when liquid or solid is at equilibrium with its gas phase ▪ Increasing temperature increases vapor pressure because vaporization is endothermic
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Effect of Volume on VP
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Increase Volume.. ▪ Pressure decreases ▪ Rate of condensation decreases
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Boiling Point
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▪ Increases as strength of intermolecular forces increase ▪ Normal boiling points = temperature at which vapor pressure of liquid = 1 atm ▪ Boiling points of molecules with hydrogen bonding are much higher than expected
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Molar heat of fusion (ΔHfus)
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▪ Heat absorbed by one mole of solid when it melts to give liquid at constantT and P
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Molar heat of vaporization (ΔHvap )
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▪ Heat absorbed when one mole of liquid is changed to one mole of vapor at constant T and P
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Molar heat of sublimation (ΔHsub )
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▪ Heat absorbed by one mole of solid when it sublimes to give one mole of vapor at constant T and P
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Energies of Phase Changes
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▪ Expressed per mole ▪ All of these quantities tend to increase with increasing intermolecular forces ▪ Molar heat of fusion (ΔHfus) ▪ Molar heat of vaporization (ΔHvap ) ▪ Molar heat of sublimation (ΔHsub )
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Phase Diagrams
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▪ Show the effects of both pressure and temperature on phase changes ▪ Boundaries between phases indicate equilibrium
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Triple point
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▪ The temperature and pressure at which s, l, and g are all at equilibrium
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Critical point
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▪ The temperature and pressure at which a gas can no longer be condensed ▪ TC = temperature at critical point ▪ PC = pressure at critical point
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Supercritical Fluid
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▪ Substance with temperature above its critical temperature (TC) and density near its liquid density ▪ Have unique properties that make them excellent solvents ▪ Values of TC tend to increase with increased intermolecular attractions between particles
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Le Chatelier’s Principle
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▪ Equilibria are often disturbed or upset ▪ When dynamic equilibrium of system is upset by a disturbance ▪ System responds in direction that tends to counteract disturbance and, if possible, restore equilibrium
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Position of equilibrium
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Used to refer to relative amounts of substance on each side of double (equilibrium) arrows
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Liquid Vapor Equilibrium
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▪ Increasing T = ▪ Increases amount of vapor ▪ Decreases amount of liquid
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Crystalline Solids
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▪ Solids with highly regular arrangements of components
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Amorphous Solids
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▪ Solids with considerable disorder in their structures
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Lattice
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▪ Many repeats of unit cell ▪ Regular, highly, symmetrical system ▪ Three (3) dimensional system of points designating positions of components (atoms, ions, molecules)
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Unit Cell
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▪ Smallest segment that repeats regularly ▪ Smallest repeating unit of lattice
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Simple cubic
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▪ Has one host atom at each corner ▪ Edge length a = 2r ▪ Where r is radius of atom or ion
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▪ Face-centered cubic (FCC)
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▪ Has one atom centered in each face, and one at each corner
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Body-centered cubic (BCC)
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▪ Has one atom at each corner and one in center
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Ionic solids lattices of alternating charges
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▪ Want cations next to anions ▪ Maximizes electrostatic attractive forces ▪ Minimizes electrostatic repulsions
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Ionic solids based on one of three basic lattices:
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▪ Simple cubic ▪ Face centered cubic ▪ Body centered cubic
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Four types of sites in unit cell
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▪ Central or body position ▪ Face site ▪ Edge site ▪ Corner site
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Central or body position
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atom is completely contained in one unit cell
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Face site
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atom on face shared by two unit cells
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Edge site
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atom on edge shared by four unit cells
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Corner site
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atom on corner shared by eight unit cells
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Factors Affecting Crystalline Structure
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▪ Size of atoms or ions involved ▪ Stoichiometry of salt ▪ Materials involved ▪ Some substances do not form crystalline solids
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Ionic Crystals
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▪ Have cations and anions at lattice sites ▪ Are relatively hard ▪ Have high melting points ▪ Are brittle ▪ Have strong attractive forces between ions ▪ Do not conduct electricity in their solid states ▪ Conduct electricity well when molten ▪ EX: NaCl, NaNO3
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Covalent Crystals
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▪ Lattice positions occupied by atoms that are covalently bonded to other atoms at neighboring lattice sites ▪ Also called network solids ▪ Interlocking network of covalent bonds extending all directions
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Covalent crystals tend to
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▪ Be very hard ▪ Have very high melting points ▪ Have strong attractions between covalently bonded atoms
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Metallic Crystals
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▪ Conduct heat and electricity ▪ Have the luster characteristically associated with metals

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