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Ch. 6 The Structure of Matter

Ch. 6 The Structure of Matter. The Importance of BONDING. Important Terms. Element = a pure substance that cannot be separated or broken down into simpler substances by chemical means Atom = the smallest unit of an element that maintains the chemical properties of that element

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Ch. 6 The Structure of Matter

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  1. Ch. 6The Structure of Matter The Importance of BONDING

  2. Important Terms • Element = a pure substance that cannot be separated or broken down into simpler substances by chemical means • Atom = the smallest unit of an element that maintains the chemical properties of that element • Compound = a substance made up of atoms of two or more different elements joined by chemical bonds

  3. Bonding • Atoms with unfilled valence shells are considered unstable. • Atoms will try to fill their outer shells by bonding with other atoms. • Chemical bond = the attractive force that holds atoms or ions together in a compound

  4. Atomic Bonds • Atoms form atomic bonds to become more stable. • Atoms become more stable by filling their valence shell or at least meeting the Octet rule by getting 8 valence electrons. Exception to Octet Rule of 8 valence electrons: Helium—which only has 1 energy level and holds a max. of 2 electrons

  5. Atomic Bonds • There are three main types of chemical bonds used by atoms to fill their valence shell: • Covalent • Metallic • Ionic “Bond, Chemical Bond”

  6. Chemical Formulas • A chemical formula tells us: • the type of atoms present • the number of atoms present • the type of compound

  7. Chemical Formulas • Example: table salt: Sodium Chloride • Chemical formula: • NaCl • Count the atoms present: • 1 Na atom • 1 Cl atom

  8. Chemical Formulas A subscript is a small number that is in a chemical formula. If no subscript is present assume that it is 1. • Example - water: H2O • 2 H atoms • 1 O atom Subscript

  9. Chemical Formulas • Sometimes there are parentheses with a subscript. The subscript only applies to the atoms within the parentheses. • Example - calcium hydroxide (kidney stones): Ca(OH)2. Ca = 1 atom O = 2 atoms H = 2 atoms

  10. Chemical Formulas • Sometimes there are subscripts in the parentheses. Multiply the subscript outside the parentheses by the subscript of each element within the parentheses. • Example - calcium nitrate: Ca(NO3)2 • 1 Ca atom • 2 N atoms • 6 O atoms (3 oxygens x 2 = 6)

  11. Covalent Bonds • Covalent bonds form between two non-metals. Groups 14-17 on the Periodic Table (plus Hydrogen) • Covalent bonds are formed when atoms SHARE electrons. • Both atoms need to gain electrons to become stable, so they share the electrons they have. • Atoms can share more than one pair of electrons to create double and triple bonds.

  12. Properties of Covalent Compounds • Results in a NEUTRAL molecule • Weak bonds • Physical State usually liquids or gases • Low Melting and Boiling Points • Poor conductors of electricity (no free electrons to move around)

  13. Covalent Bonds Use Lewis structures to draw valence electrons for each atom in the covalent pair. Cl Cl Each chlorine atom wants to gain one electron to achieve an octet.

  14. Covalent Bonds The octet is achieved by each atom sharing the electron pair in the middle. Cl Cl octet octet Now, each Chlorine atom has 8 valence electrons because it is sharing one pair.

  15. Chlorine Molecule It is a single bonding pair so it is called a single covalent bond. The compound is now called a molecule. Cl Cl Cl Cl Cl2

  16. Covalent Bonds O O How will oxygen bond?

  17. Covalent Bonds Two bonding pairs, making a double bond. O O The double bond can be shown as two dashes OO O2

  18. Covalent Bonds • Elements can share up to three pairs of electrons. (6 total electrons). Single Bond (2e) Double Bond (4e) Triple Bond (6e)

  19. Covalent Bonds • Atoms can share their electrons equally or unequally. • When atoms share electrons equally, it is called a non-polar covalent bond. • Non-polar covalent bonds form between atoms of the same type. Ex: H2, Cl2, • When atoms share electrons unequally it is called a polar covalent bond. • One atom pulls the electrons closer to itself. • The atom that pulls the electrons more gets a slightly negative charge. • The other atom gets a slightly positive charge. • Ex: Water molecule [*Animation @ 0:54 sec for covalent] Bonding Animation

  20. Covalent Bonds Nomenclature • Naming binary covalent compounds: • Two nonmetals • Name each element • Change the ending of the 2nd element to –ide • Use prefixes to indicate the # of atoms of each element • Do not use “mono” with the first element

  21. Covalent Bonds Nomenclature • CO • carbon monoxide • CO2 • carbon dioxide • PCl3 • phosphorus trichloride • CCl4 • carbon tetrachloride • N2O • dinitrogen monoxide

  22. Covalent Bonds Nomenclature Given the following covalent compounds, WRITE the correct chemical formula. HS2 P2O5 N3F6

  23. Practice: Drawing Covalent Bonds • We can illustrate covalent bonding using Lewis structures. • 1 – Draw a Lewis structure for each element. • Ex: C H • 2 - Continue adding atoms until all atoms have a full valence H H C H CH4 carbon tetrahydride H

  24. Ions • Ions are formed when atoms gain or lose electrons. • Ions are charged atoms (positive or negative). • Positive ions are called cations. • Formed when the atom loses electrons. • Lose negative charge, becomes positive ION • Metals • Negative ions are call anions. • Formed when the atom gains electrons. • Gain negative charge, become negative ION • Non-metals

  25. Ionic Bonds • Ionic bonds are formed between metals and non-metals. • Ionic bonds are formed between oppositely charged atoms (ions). • Ionic bonds are formed by the transfer of electrons. • One atom loses (gives away) electrons. • One atom gains (receives) electrons.

  26. Ionic Bonds • Use the number of valence electrons to determine the # of electrons that are lost or needing to be gained. • The transfer of electrons create a positive ion and a negative ion. The opposite charges attract one another, causing a chemical bond to form. Bonding Animation

  27. Atoms with 4 or less valence electrons want to LOSE (give away) their valence electrons. [Groups 1, 2, 13, 14] • Atoms with 4 or more valence electrons want to GAIN (receive) more electrons to satisfy their octet. [Groups 14, 15, 16, 17]

  28. Ionic Bonds • The normal charge of an ion can be quickly determined using the oxidation number of an element. • The oxidation numberof an atom is the charge that atom would have if the compound was composed of ions.

  29. Ionic Bonds • To find the oxidation number : • Look at Group # • Determine # of valence electrons • If 4 or less, atom will lose (give away) valence electrons (ion is positive) • If 4 or more, atom will gain the needed # to fill valence shell. (ion is negative)

  30. Ionic Bonds • Example: • Beryllium is in Group 2 • Be has 2 e- • Wants to achieve octet • Loses the 2 e- • Oxidation #/Ion charge of +2 • Example: • Nitrogen is in Group 15 • N has 5 e- • Needs 3 more for octet • Gains 3 e- • Oxidation #/Ion charge of -3

  31. Practice: Determining Oxidation Numbers 16 6 -2 +2 2 2 -1 17 7 15 5 -3 1 1 +1

  32. Ionic Bonding Nomenclature To name Binary Ionic Compounds: • 2 elements—one METAL and one NON-METAL • Cation is always written first [Metal] • Cation name stays the same • Anion is written second [Non-metal] • Change the non-metal’s ending to “-ide”. • NO PREFIXES ARE USED FOR IONIC COMPOUND NAMING

  33. Examples NaCl Name the metal ion Sodium Chloride CaO Name the nonmetal ion, changing the suffix to –ide. Oxide Calcium Al2S3 Aluminum Sulfide MgI2 Magnesium Iodide This is two metals – not a binary ionic compound BaNa2 The name of this is Banana (JOKE – haha) You should recognize a problem with this one

  34. Drawing Ionic Bonds • 1 – Draw the Lewis structure for each element. • Ex: Na Cl • 2 – Draw arrows to show the TRANSFER (gain/loss) of electrons [draw extra atoms if needed]

  35. Drawing Ionic Bonds (continued) • 3 – Draw ion Lewis diagrams showing the new charge for each ion. • Ex: • 4- Write the chemical formula for the compound formed represents the ratio of negative ions to positive ions. • Ex: NaCl – for every 1 sodium ion, there is also 1 chlorine ion. Chemical Formula = NaCl

  36. Practice Drawing Ionic Bonds Elements Lewis Transfer Formula Diagram Calcium Fluorine Sodium Oxygen

  37. “Swap & Drop” Method Given the name of an Ionic Compound, you can determine the chemical formula using the “swap and drop” method: Write the symbols for each ion. Determine the oxidation number of each ion. Swap and Drop Reduce (if necessary). Rewrite Be F 2+ 1- beryllium fluoride BeF2

  38. Ionic Bonds Form when electrons are transferred between atoms. Form between a metal and a non-metal. Covalent Bonds Form when electrons are shared between atoms. Form between twonon-metals. Ionic vs. Covalent Bonds in Binary Compounds Both types of bonds result in all atoms having a full outer energy level.

  39. Ionic vs. Covalent Bonds in Binary Compounds Other comparisons between Ionic and Covalent Compounds: Ionic Compounds • Results in a Neutral Compound • Crystalline Solid • Strong Bonds • High Melting Point Covalent Compounds • Results in a Neutral Molecule • Mostly results in gases or liquids • Weak Bonds • Low Melting Points

  40. Dogs teaching Chemistry https://www.youtube.com/watch?v=_M9khs87xQ8&edufilter=ITEV1GgfY7rUuGxSyy0SA

  41. Polyatomic Ions • A polyatomic ion is a group of covalently bonded atoms that have lost or gained an electron. (Example: Nitrate NO3- and Ammonium NH4+). • Oppositely charged polyatomic ions can form compounds. (Example: Ammonium nitrate NH4NO3).

  42. Polyatomic Ions • Naming of these compounds follows the same rules as binary ionic compounds. • The most important part is recognizing there is a polyatomic ion present.

  43. Practice: Polyatomic Ions To go from the formula to the name: • Name the cation. • Name the anion. Be(NO3)2 (NH4)2S beryllium ammonium nitrate sulfide

  44. Polyatomic Ions To go from name to formula: Write the symbols for each ion. Determine the oxidation number of each ion. Swap and Drop Reduce (if necessary). Put parentheses around the polyatomic ion if receives a subscript greater than one. Rewrite ammonium oxide NH41+ O2- (NH4)2O ** Remember charges CANCEL out each other!!

  45. Practice: Polyatomic Ions Ca2+ PO43- Ca3(PO4)2 Na1+ OH1- NaOH (NH4)2SO4 NH41+ SO42-

  46. Metallic Bonds • Metallic bonds are metal to metal bonds formed by the attraction between positively charged metal ions and the electrons around them. • Atoms are packed tightly together to the point where outermost energy levels overlap. • This allows electrons to move freely from one atom to the next making them great conductors of electricity.

  47. Transition Metals--Ionic Compounds • Transition metals are cations that have variable charges that makes them hard to name. • We use Roman numerals to indicate the charge of a transition metal. • Example: • copper (II) oxide – charge of copper for this compound is +2 • titanium (IV) sulfide – charge of titanium for this compound is +4

  48. Transition Metal Ionic Compounds 2+ 2- • To go from formula to name you need to determine the Roman numeral for your transition metal. • If there are no subscripts, simply give the transition metal the equal and opposite charge to the nonmetal. • Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY. FeO iron (II) oxide

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