In chemistry, a bond or chemical bond is a link between atoms in molecules or compounds and between ions and molecules in crystals. A bond represents a lasting attraction between different atoms, molecules, or ions.
1. Ch. 6 The Structure of Matter The Importance of BONDING
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5. 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
6. 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
7. Atomic Bonds • Atoms form atomic bonds to become more stable. – Atoms become more stable by filling their valence shell or at Exception to Octet Rule least meeting the octet of 8 valence electrons: rule by getting 8 Helium—which only has valence electrons. 1 energy level and holds a max. of 2 electrons
8. Atomic Bonds • There are three main types of chemical bonds used by atoms to fill their valence shell: “Bond, – Covalent Chemical – Metallic Bond” – Ionic
9. Chemical Formulas • A chemical formula tells us: – the type of atoms present – the number of atoms present – the type of compound
10. Chemical Formulas • Example: table salt: Sodium Chloride • Chemical formula: – NaCl • Count the atoms present: – 1 Na atom – 1 Cl atom
11. Chemical Formulas • Sometimes there are subscripts present. – 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
12. 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. – 1 Ca atom – 2 O atoms – 2 H atoms
13. 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)
14. Covalent Bonds • Covalent bonds form between two non- metals. Groups 14-17 on the Periodic Table • 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.
15. Properties of Covalent Compounds Results in a NEUTRAL molecule Weak bonds Physical State usually liquids or Low Melting and Boiling Points Poor conductors of electricity (no free electrons to move around)
16. Covalent Bonds Use Lewis structures to draw valence electrons for each atom in the covalent pair. Each chlorine atom wants to gain one electron to achieve an octet.
17. Covalent Bonds The octet is achieved by each atom sharing the electron pair in the middle. Now, each Chlorine atom has 8 valence electrons because it is sharing one pair.
18. 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 Cl2
19. Covalent Bonds How will oxygen bond?
20. Covalent Bonds Two bonding pairs, making a double bond. The double bond can be shown as two dashes O O O2
21. Covalent Bonds • Elements can share up to three pairs of electrons. (6 total electrons). Single Bond (2e) Double Bond (4e) Triple Bond (6e)
22. 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 Bonding Animation
23. Covalent Bonds Nomenclature • Naming binary covalent compounds: # of Atoms Prefix – Two nonmetals 1 mono- – Name each element 2 di- – Change the ending of 3 tri- the 2nd element to 4 tetra- –ide 5 6 penta- hexa- – Use prefixes to 7 hepta- indicate the # of atoms of 8 octa- each element 9 nona- – Do not use “mono” with the first element 10 deca-
25. Covalent Bonds Nomenclature Given the following covalent compounds, WRITE the correct chemical formula. Name Chemical Formula Hydrogen Disulfide HS2 Diphosphorus pentoxide P2 O5 Trinitrogen hexafluoride N3F6
26. 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 carbon tetrahydride H
27. 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
28. 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.
29. 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
30. 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]
31. Ionic Bonds • The normal charge of an ion can be quickly determined using the oxidation number of an element. – The oxidation number of an atom is the charge that atom would have if the compound was composed of ions.
32. 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)
33. 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
35. 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
36. Sodium Chloride Name the metal ion Calcium Oxide Name the nonmetal Al2S3 ion, changing the Aluminum Sulfide suffix to –ide. Magnesium Iodide BaNa2 Thisshould The You is two of name metals this is–Banana recognize not a binary (JOKE a problem ionic – this with haha)one compound
37. 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]
38. 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
39. Practice Drawing Ionic Bonds Elements Lewis Transfer Formula Diagram
40. “Swap & Drop” Method Given the name of an Ionic Compound, you can determine the chemical formula using the “swap and drop” method: 1. Write the symbols for each ion. 2. Determine the oxidation number of each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. Rewrite
41. Ionic vs. Covalent Bonds in Binary Compounds Ionic Bonds Covalent Bonds • Form when electrons • Form when electrons are transferred are shared between between atoms. atoms. • Form between a • Form between two metal and a non- non-metals. metal. Both types of bonds result in all atoms having a full outer energy level.
42. Ionic vs. Covalent Bonds in Binary Compounds Other comparisons between Ionic and Covalent Compounds: Ionic Compounds Covalent Compounds • Results in a • Results in a Neutral Neutral Compound Molecule • Crystalline Solid • Mostly results in • Strong Bonds gases or liquids • High Melting • Weak Bonds Point • Low Melting Points
43. 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).
44. Polyatomic Ions • Naming of these Common Polyatomic Ions ammonium NH4+ compounds follows carbonate CO32- the same rules as bicarbonate HCO3- binary ionic hydroxide OH- compounds. nitrate NO3- – The most important nitrite NO2- part is recognizing phosphate PO43- there is a polyatomic sulfate SO42- ion present. sulfite SO32- acetate C2H3O2-
45. Practice: Polyatomic Ions To go from the formula to the name: 1. Name the cation. 2. Name the anion.
46. Polyatomic Ions To go from name to formula: 1. Write the symbols for each ion. 2. Determine the oxidation number of O 2- each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. Put parentheses around the polyatomic (NH4)2O ion if receives a ** Remember charges CANCEL subscript greater than out each other!! one. 6. Rewrite
47. Practice: Polyatomic Ions Compound Name Oxidation #s Chemical Formula Calcium phosphate Ca2+ PO43- Ca3(PO4)2 Sodium hydroxide Na1+ OH1- NaOH Ammonium sulfate (NH4)2SO4 NH41+ SO42-
48. 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.
49. 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
50. Transition Metal Ionic Compounds • To go from formula to name you need to determine the Roman numeral for your transition metal. 1. If there are no subscripts, simply give the transition metal the equal and opposite charge to the nonmetal. 2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.
51. Transition Metal Ionic Compounds • To go from formula to name you need to determine the Roman numeral for your transition metal: 1. If there are subscripts present use the reverse “Swap and Drop.” 2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.
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