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Chemical Bonds

Chemical Bonds

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Chemical Bonds

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  1. Chemical Bonds

  2. 6-1 Ionic Bonding • When is an atom unlikely to react? • What is one way in which elements can achieve stable electron configurations? • How does the structure of an ionic compound affect its properties?

  3. Stable Electron Configurations • When the highest occupied energy level of an atom is filled with electrons, the atom is stable and not likely to react. • The highest occupied energy level of a noble gas atom is filled. • The noble gases have stable electron configurations with eight valence electrons (or 2 in the case of helium)

  4. The Electron Dot Diagram • Since the chemical properties of an element depend on the number of valence electrons, it is useful to have a model that focuses on valence electrons. • A model of an atom in which dot represents valence electrons • Symbol in the center represents the nucleus and all the other electrons

  5. Ionic Bonds • Elements that do not have complete sets of valence electrons tend to react. • By reacting, they achieve electron configurations similar to those of noble gases. • Some elements achieve stable electron configurations through the transfer of electrons between atoms. • Transfer of Electrons Chlorine to Argon Sodium to Neon

  6. Ionic Bonds Continued • Formation of Ions • When atoms gain or lose an electron, the number of protons is no longer equal to the number of electrons. • The charge is not balanced and atom is not neutral. • An atom that has a net positive or negative electric charge is an ion. This is represented by a + or – sign next to the symbol. • Chlorine gains an electron =17 protons, 18 electrons = 1- charge • Sodium has 11 protons and 10 electrons (one extra proton) = 1+ charge Negative charge - anion…suffix = ide Positive charge – cation… just use element name

  7. Ionic Bonds Continued • Formation of Ionic Bonds • Particle with negative charge will attract a particle with a positive charge • When anion and cation are close, a chemical bond forms. • Chemical bond: force that holds atoms or ions together as a unit (attraction between them) • Ionic bond: force that holds cations and anions together (when electrons are transferred from one atom to another) • Ionization Energy: • The amount of energy used to remove an electron • An electron can move to a higher energy level when an atom absorbs energy because the energy allows the electron to overcome the attraction of the protons in the nucleus.. • Varies from element to element • The lower the energy, the easier it is to remove an electron from an atom

  8. Ionic Compounds • Compounds that contain ionic bonds are ionic compounds. • can be represented by chemical formulas (a notation that shows what elements a compound contains and the ratio of the atoms or ions of those elements in the compound) • Example: NaCl = one sodium ion for each chloride ion in sodium chloride • Magnesium chloride = ?

  9. Ionic Compounds Continued • To review • A chemical formula for an ionic compound tells you the ratio of the ions in the compound • It does not tell you how the ions are arranged in the compound • Solids arranged in a lattice structure are called crystals

  10. Ionic Compounds Continued • A: sodium chloride: ions arranged in an orderly three-dimensional structure (cubic shape) Crystals: solids arranged in a lattice structure • B: sodium chloride: crystals look like cubes The properties of an ionic compound can be explained by the strong attractions among ions within a crystal lattice. Class Participation Opportunity: Research report on how to make rubies

  11. 6-2 Covalent Bonding • How are atoms held together in a covalent bond? • What happens when atoms don’t share electrons equally? • What factors determine whether a molecule is polar? • How do attractions between polar molecules compare to attractions between nonpolar molecules?

  12. Covalent Bonds • Plants absorb water through roots • Carbon dioxide from the air enters through stomata in leaves • Plants use energy from sun to convert water and carbon dioxide into sugar • Energy is stored in the chemical bonds of the sugar • The elements of sugarare carbon, oxygen and hydrogen(nonmetals with high ionization energies so a transfer of electrons does not tend to occur). • When nonmetals join, they share valence electrons forming a covalent bond.

  13. Covalent Bonds Continued • The attractions between the shared electrons and the protons in each nucleus hold the atoms together in a covalent bond. • Sharing Electrons: • A hydrogen atom has one electron • Two hydrogen atoms can achieve a stable electron configuration by sharing their electrons • When two atoms share one pair of electrons the bond is called a single bond

  14. Covalent Bonds Continued • Four different ways to represent a covalent bond. • Electron dot model = pair of dots • Structural formula = line • Electron cloud model and space-filling model = orbitals of atoms overlap • Hydrogen atoms bonded together form a molecule (a neutral group of atoms joined together by one or more covalent bonds). • The attraction between the share electrons and the protons in each nucleus hold the atoms together in a covalent bond.

  15. Covalent Bonds Continued Many nonmetal elements exist as diatomic molecules (two atoms). • Why are the atoms in the models of diatomic molecules not complete spheres? • Why are the spheres in the models of fluorine, chlorine, and bromine different sizes? The space-filling models show that orbitals of atoms overlap when they form covalent bonds. The different sizes of spheres model the different atomic radii of the atoms.

  16. Covalent Bonds Continued • Multiple Covalent Bonds • Nitrogen = five valence electrons • Two nitrogen atoms = shared a pair of electrons, each would have seven • Nitrogen molecule = each molecule shared three pairs of electrons, each atoms has eight valence electrons • triple bond

  17. Unequal Sharing of Electrons Generally, elements on the right and on the top of the periodic table have a greater attraction for electrons that elements on the left (except noble gases). • In Polar Covalent Bonds, electrons are not shared equally. • When atoms form a polar covalent bond, the atom with the greater attraction for electrons has a partial negative charge. The other atom has a partial positive charge. • Polar bonds occur in molecules have only two atoms.

  18. Unequal Sharing of Electrons Continued • Polar and Nonpolar Molecules • The type of atoms in a molecule and its shape are factors that determine whether a molecule is polar or nonpolar. • In carbon dioxide, there are double bonds between each oxygen atom and the central carbon atom (because oxygen has a greater attraction for electrons than carbon does, each double bond is polar). The molecule is linear. Oxygen atoms are opposite each other. There is an equal pull on the electrons from opposite directions. These pulls cancel out and the molecule is nonpolar. • In a water molecule, there are two single polar bonds. Oxygen has a greater attraction for electrons than hydrogen does. The molecule is bent; the polar bonds do not cancel out. Oxygen has partial negative charge; hydrogen has a partial positive charge.

  19. Attraction Between Molecules There are forces of attraction between molecules but they are not as strong as the attractions between ionic and covalent bonds. They are strong enough to hold the molecules together in a liquid or solid. • Attractions between polar molecules are stronger than attractions between nonpolar molecules.

  20. 6-3 Naming Compounds and Writing Formulas • What information do the name and formula of an ionic compound provide? • What information do the name and formula of a molecular compound provide?

  21. Chemists name compounds based on their compositions to avoid confusion. • Thomas Drummond discovered the white solid lime(calcium oxide) that emits a bright light when heated to a high temperature. • Lime wash • Quicklime • Unslaked lime • To avoid confusion, chemists call lime calcium oxide. Can be mixed with paint for whitewashing

  22. Describing Ionic Compounds The vase and the plate are both coated in oxides of copper (copper and oxygen), one red and one black. Calling both oxides by the same name won’t work, so there must be two names. • The name of an ionic compound must distinguish the compound from other ionic compounds containing the same elements. • The formula of an ionic compound describes the ratio of the ions in the compound.

  23. Describing Ionic Compounds Continued Binary Ionic Compounds A compound made from two elements is a binary compound. The name of the cation is followed by the name of the anion. cation = metal without any change sodium atom and sodium ion anion = uses part of the name of the nonmetal with the suffix ide iodine atom and iodide ion

  24. Describing Ionic Compounds Continued • Metals with Multiple Ions • Many transition metals form more than on type of ion • Copper (I) 1+ • Copper (II) 2+ • Red copper (I)oxide – “copper one oxide” Cu2O →Cu1+O2- • Black copper (II) oxide – “copper two oxide” CuO→ Cu2+ + O2-

  25. Describing Ionic Compounds Continued • Polyatomic Ions – ammonium ion in picture (below right) • A covalently bonded group of atoms that has a positive or negative charge on and acts as a unit is a polyatomic ion. • Most are anions (below left) • In iron (III) hydroxide, the subscript 3 indicates that there are three hydroxide ions for each iron (III) ion.

  26. Writing Formulas for Ionic Compounds • Write the cation first, followed by the symbol of the anion • Use subscripts to show the ratio of the ions in the compound • Since all compounds are neutral, the total charges on the cations and anions must add up to zero Two Na (Na+) for each S(S2-) Na2S

  27. Writing a Formula for Ionic Compounds

  28. Describing Molecular Compounds • The name and formula of a molecular compound describe the type and number of atoms in a molecule of the compound(unlike ionic compounds). • In other words, the names and formulas identify (match) specific compounds. • Most metallic elements appears first in the name (farther to the left in periodic table or if in same group, closer one to the bottom) • Second element is changed t end in the suffix –ide • Examples: • carbon dioxide - CO2 • Dinitrogentetraoxide - N2O4 • Mononitrogen dioxide, commonly written as nitrogen dioxide - NO2 • Diphosphorustetraflouride - P2F4

  29. 6-4 The Structure of Metals • What are the forces that give a metal its structure as a solid? • How do metallic bonds produce some of the typical properties of metals? • How are the properties of alloys controlled?

  30. The Light Bulb: A New Technology? • Just before the year1900, the light bulb was cutting edge technology. • Researchers were busy working on finding the best material for the filaments • *had to be ductile enough *couldn’t melt • *had to have low vapor pressure • Found Tungsten (W) • *highest melting point • *lowest vapor pressure

  31. Metallic Bonds • Metal atoms achieve stability by losing electrons. • If there are no nonmetal atoms available to accept the electrons, valence electrons move among the atoms. (The metal atoms become cations surrounded by a pool of shared electrons.) • This attraction between metal cation and the shared electrons is a metallic bond. • The cations in a metal form a lattice that is held in place by strong metallic bonds between the cations and the surrounding valence electrons.

  32. Metallic Bonds Continued Even though the electrons are moving among the atoms, the total number of electrons does not change so the metal is neutral. The more valence electrons an atom can contribute to the pool, the stronger the bond will be. alkali metal = weak bond (1 valence electron), soft with low melting point transition metal = stronger bond )more valence electrons), harder with higher melting points

  33. Explaining Properties of Metals • The structure within a metal affects its properties. • The mobility of electrons within a metal lattice explains some of the properties of metals. • Examples: ability to conduct an electric current and malleability • Conducting electricity: • Metals have a pool of electrons, a built-in supply of charged particles that can flow from one location to another • Electric currents are carried by these free flowing shared electrons. • Malleability • The lattice in metals is flexible (unlike in rigid ionic compounds). • This is why tungsten and copper can be made into thin wires without breaking. • Picture below illustrates how ions shift in position and shape changes but the metal does not shatter because of metallic bonds between the ions and electrons.

  34. Alloys • Scientists can design alloys with specific properties by varying the types and amounts of elements in an alloy. • Gold • 100% pure gold is expressed in karats…24 karat gold • Soft metal that can be easily worn away • Gold alloys • Gold mixed with silver, copper, nickel or zinc gold is more resistant to wear • 12 karat gold is only 50% gold • 18 karat gold is 75% gold

  35. Copper and Steel Alloys • Copper Alloys • First most important alloy of copper = bronze • Associated with the Bronze Age • In its simplest form, bronze contains only copper and tin, which when mixed together, are stronger than they are alone • Can be used for ships, statues and bells • Brass is another alloy of copper (contains copper and zinc) • Bronze and brass have distinctly different properties • Brass shiner, softer, weathers more quickly • Steel Alloys • 1900s = Steel Age = skyscrapers, automobiles and ships • Alloy of iron and some carbon • Properties depend on an element other than iron and carbon • Stainless steel = more than 10% chromium = durable and resists rust • Steel cables = sulfur, manganese phosphorus and silicon = stronger

  36. Another Example • Airplane parts • Aluminum is lightweight but too soft and dents too easily • With a small amount of copper or manganese, the aluminum is stronger but still lighter than steel

  37. A Historical Perspective