CHEMICAL BONDING [No one wants to be alone] The Marrying of Atoms (AIM)

Transcription

1 CHEMICAL BONDING [No one wants to be alone] The Marrying of Atoms (AIM) Associate Degree in Engineering Prepared by M. J. McNeil, MPhil. Department of Pure and Applied Sciences Portmore Community College Main Campus

2 LECTURE OBJECTIVES Explain the formation of ionic and covalent bonds and draw dot and cross diagram to illustrate the bonding. Predict the likelihood of an atom to forming an ionic or covalent bond based on atomic structure. Explain metallic bonding (arrangement of cations and mobile electrons).

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4 PROPERTIES OF MATTER Macroscopic properties of matter vary greatly due to the type of bonding.

5 WHAT IS CHEMICAL BONDING? How could you get pieces of paper to bond (stick together)? Chemical bonds is the chemical combination (sticking together) of atoms. Can form by The attraction of positive ion to a negative ion or The attraction of the positive nucleus of one atom and the negative electrons of another atom. Atoms combine by forming bonds, when they do so new types of particles are formed compounds.

6 BONDING OF ATOMS (KEYS) AS ATOMS THEY CANNOT ATTRACT!!! Bonding involves Transfer of electrons: Lose electrons Gain electrons Chemical compounds are formed by the joining of two or more atoms. A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. The bound state implies a net attractive force between the atoms... a chemical bond. Sharing of electrons Valence electrons - number of unpaired (single) electrons in the outermost shell

7 WHAT HAPPENS DURING BONDING? Atoms stick together when their outer shell or valence electrons interact to form bonds. Atoms may lose or gain or share their valence electrons. As these changes occur, the electronic configurations of the atoms change and new particles are formed. These are ions or molecules). The new configuration of each atom will (or will appear to) adopt the electronic configuration of the nearest NOBEL GAS in the Periodic Table. Noble gases have filled valence shells which are stable electron configurations.

8 WHY ATOMS FORM BONDS? Atoms form bonds with one another because they want to become STABLE. Atoms can achieve this FULL OCTET of eight electrons or gaining and losing electrons: whether electrons are shared or transferred determines which type of bonding is formed. The aim of bonding is to also get the outer shell sometimes EMPTIED. There are THREE TYPES OF BONDING.

9 WHICH ATOMS LOSE, GAIN OR SHARE ELECTRONS? Metal atoms containing 1, 2 or 3 valence electrons tend to lose electrons. (The larger the atomic radius of the atom, the more easily it loses electrons) Some non-metal atoms with 5, 6, 7 valence electrons may gain or share their electrons to fill their valence shells. The smaller the non-metal atoms, the more easily it readily accepts electrons. Non-metals containing 4 to 7 valence electrons may also share electrons when combining with other non-metals. Elements having different electronic configuration. Different electronic arrangement means different types of bonding.

10 TYPES OF CHEMICAL BONDS - AN ATTEMPT TO FILL EMPTY SHELLS There are THREE MAIN types of bonding: 1. Ionic or Electrovalent Bonding 2. Covalent Bonding 3. Metallic Bonding AS ATOMS THEY CANNOT ATTRACT!!!

11 IONIC (ELECTROVALENT) BONDING Ionic Bonding is the complete transfer of electrons from one atom (usually a metal) to another atom (a non-metal) with very different electronegativity. The metal loses electrons and the non-metal gains. As a result of ionic bonding, ions are created, which are charged atoms. These oppositely charged cations and anions are attracted to one another because of their opposite charges. There is a chemical bond that is formed between oppositely charge ionic particles. Ionic

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13 IONS If it loses an electron, it become a ion with a positive charge known as CATION. (Why? Illustrate!!!) If it gains an electron, it becomes a ion with a negative charge known as an ANION) (Why? Illustrate!!!!)

14 ATOM E.C. Lose / gain electrons Charge on ion H 1 He 2 Li 2,1 Be 2,2 B 2,3 C 2,4 N 2,5 O 2,6 F 2,7 Ne 2,8

15 ATOM E.C. Lose / gain electrons Charge on ion Na 2,8,1 Mg 2,8,2 Al 2,8,3 Si 2,8,4 P 2,8,5 S 2,8,6 Cl 2,8,7 Ar 2,8,8 K 2,8,8,1 Ca 2,8,8,2

16 COPY AND COMPLETE THE TABLE p n e A (mass #) e.c. charge ,8, ,8,8

17 RULES TO DRAW DOT (. ) AND CROSS (x) DIAGRAMS These diagrams are used to illustrate the formation of bonds. The (. ) and (x) symbols are used to indicate which atom the electrons came from. (1) Determine if the compound of interest has been formed from ionic or covalent bonding. [If a metal is present, the compound is formed from ionic bonding. Otherwise its, covalent.] (2) Determine the formula of the compound using valency. If the compound is composed of ions, its is ionic in nature. (3) Draw the basic atomic structure of the element separately to show its valence electrons. Use different symbols to represent these electrons for each of the drawn atomic structures. (4) Draw arrows to showcase electrons which are transferred or electrons (ionic bonding) or which are shared (covalent bonding). (5) Redraw the ions (cation - positive charged ion; anions - negatively charged ion) after the electron transfer or the molecule after the electron sharing.

18 Ionic bond - electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na + ) and the Cl becomes (Cl - ), charged particles or ions. The positive sodium ion and the negative chloride ion are strongly attracted to each other. This attraction, which holds the ions close together, is a type of chemical bond called an ionic bond.

19 CONNECTING SODIUM AND CHLORINE Chlorine steals one of sodium s electrons

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21 Positive Ion (Cation) Formation Na has one valence electron. It loses it to Chlorine. Na now has a filled valence shell. (an octet) Becomes positive one in charge Chlorine has seven valence electrons. It gains one electron from Na. Chlorine now has filled octet. Chlorine has a negative one charge. (Chloride ion) Na +1 attracts Cl -1 and forms the ionic bond. Negative Ion (Anion) Formation

22 MAGNESIUM OXIDE The element magnesium, Mg, in Group 2 has two electrons in its outer energy level. Magnesium can lose these two electrons and achieve a completed energy level. MgO (Chemical Formula)

23 IONIC COMPOUND FORMULA

24 TRY THESE Draw the bonding that occurs in the following compounds. 1. Sodium chloride 2. magnesium fluoride 3. magnesium oxide 4. aluminum oxide 5. calcium oxide 6. Li 2 O 7. Mg 3 N 2

25 COVALENT BOND Some atoms are unlikely to lose or gain electrons because the number of electrons in their outershell makes this difficult. The alternative is SHARING, not transferred. These give rise to covalent bond formation of sharing pairs. Covalent bonds are those formed between nonmetallic elements of similar electronegativity. Shared electrons are ATTRACTED to their nuclei. They move back and forth between the outer energy levels of atoms in the covalent bond. Covalent

26 COVALENT BONDING The bond arises from the mutual attraction of 2 nuclei for the same electrons. A covalent bond is a balance of attractive and repulsive forces. 26

27 COVALENT BONDING The bond arises from the mutual attraction of 2 nuclei for the same electrons. The neutral particle that is form when atoms share electrons are called MOLECULES. (Recall AIM ) A molecule is the basic unit of a MOLECULAR COMPOUND. Also, molecules are groups of atoms that are held together by covalent bonds in a specific ratio & shape. A covalent bond is a balance of attractive and repulsive forces.

28 Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Atom Oxygen Molecule (O 2 )

29 COVALENT BONDING The bonds between oxygen and hydrogen in a water molecule are covalent bonds. There are two covalent bonds in a water molecule, between the oxygen and each of the hydrogen atoms. Each bond represents one electron.

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32 BOND AND LONE PAIRS Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. (Lewis structures) H shared or bond pair Cl Unshared or lone pair (LP) This is a LEWIS ELECTRON DOT structure. 32

33 TRY THESE 1. methane 2. nitrogen 3. chlorine 4. carbon dioxide 5. water 6. ammonia 7. hydrogen cyanide FANCY BONDING Sometimes, atoms share more than one electron. Occasionally, they can share 2 or even 3 electrons. These are called double and triple bonds.

34 Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons C would like to N would like to O would like to Gain 4 electrons Gain 3 electrons Gain 2 electrons

35 IONIC COMPOUNDS -Formed from a combination of metals and nonmetals. -Electron transfer from the cation to the anion. -Opposite charged ions attract each other. Solids at room temperature High melting points Dissolve well in water Conduct electricity only when dissolved in water; electrolytes COVALENT COMPOUNDS -Formed from a combination of nonmetals. -Electron sharing between atoms. Can be solid, liquid, or gas at room temperature. Low melting points Do not dissolve in water (Sugar is an exception) Do not conduct electricity; non electrolytes Brittle, hard Soft

36 LEADING QUESTIONS Why do some solids dissolve in water but others do not? Why are some substances gases at room temperature, but others are liquid or solid? What gives metals the ability to conduct electricity, what makes non-metals brittle? The answers have to do with.

37 INTERMOLECULAR FORCES (IMF) OF ATTRACTION

38 INTERMOLECULAR FORCES (IMF) Intermolecular forces are attractive forces between molecules. VS Intramolecular forces hold atoms within a molecule. Intermolecular vs Intramolecular 41 kj to vaporize 1 mole of water (inter) 930 kj to break all O-H bonds in 1 mole of water (intra) Generally, intermolecular forces are much weaker than intramolecular forces. Measure of intermolecular force boiling point melting point etc/

39 What determines if a substance is a solid, liquid, or gas? The IMF and the temperature (kinetic energy) of the molecules. Gases: The average kinetic energy of the gas molecules is much larger than the average energy of the attractions between them. Liquids: the intermolecular attractive forces are strong enough to hold the molecules close together, but without much order. Solids: the intermolecular attractive forces are strong enough to lock molecules in place (high order). Are they temperature dependent? 39

40 TYPES OF IMFs (BETWEEN NEUTRAL MOLECULES) IMF forces are weaker than ionic, covalent and metallic bonds and there exist three main types: 1. Temporary dipole-dipole Forces (London dispersion forces/van der Waals forces) 2. Permanent dipole-permanent Dipole Forces 3. Hydrogen Bonding

41 POLARITY - TUG OF WAR

42 Van der Waal FORCES (Non-Polar Molecules) Van der Waals forces also known as London dispersion forces named after Fritz London. Non-polar molecules DO NOT have dipoles like polar molecules. How then can non-polar compounds form solids or liquids? London forces are due to small dipoles that exist in non-polar molecules. Because electrons are moving around in atoms, there will be instants when the charge around an atom is unequal. (e - equally shared) electrons are shifted to overload one side of an atom or molecule. The resulting TINY DIPOLES cause attractions between atoms/molecules.

43 Van der Waal FORCES (Non-Polar Molecules) Instantaneous dipole: Induced dipole: Eventually electrons are situated so that tiny dipoles form A dipole forms in one atom or molecule,inducing a dipole in the othe

44 Halogen Boiling Pt (K) Noble Gas Boiling Pt (K) F He 4.6 Cl Ne 27.3 Br Ar 87.5 I Kr Explain the trend in boiling points? Hint: look at the relative sizes. 44

45 Polar Covalent Bonds: Unevenly matched, but willing to share.

46 PERMANENT-DIPOLE PERMANENT DIPOLE IMF Molecules can have a separation of charge. + This happens in both ionic and polar bonds (the greater the electronegativity (EN), the greater the dipoles). Tend to exist in non-polar covalent substances where electrons are not shared equally. H Cl Molecules are attracted to each other in a compound by these +ve and -ve forces Polar molecules have dipole-dipole attractions for one another. + HCl HCl - dipole-dipole attraction

47 PERMANENT DIPOLE PERMANENT DIPOLE IMF 16 kj/mol (to separate molecules) kj/mol (to break bond) 47

48 - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen. How is structure related to chemical and physical properties?

49 POLAR AND NONPOLAR COVALENT BONDS KEYS POLAR DIFFERENT TYPES OF ATOMS DO NOT SHARE ELECTRONS EQUALLY NONPOLAR ATOMS OF THE SAME TYPE SHARE THE ELECTRONS EQUALLY

50 HYROGEN BONDING IMF (Strongest) Hydrogen bonding is a weak to moderate attractive force that exists between a hydrogen atom covalently bonded to a very small and highly electronegative atom and a lone pair of electrons on another small, electronegative atom (F, O, or N).

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52 HYDROGEN BONDING The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative F, O or N atom. A H B or A H A A & B are F, O, N

53 HYDROGEN BOND + H-F H-F -

54 TESTING CONCEPTS 1. Which attractions are stronger: intermolecular or intramolecular? 2. How many times stronger is a covalent bond compared to a dipoledipole attraction? 3. What evidence is there that nonpolar molecules attract each other? 4. Suggest some ways that the dipoles in London forces are different from the dipoles in dipole-dipole attractions. 5. A) Which would have a lower boiling point: O 2 or F 2? Explain. B) Which would have a lower boiling point: NO or O 2? Explain.

55 What type(s) of intermolecular forces exist between each of the following molecules? HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH 4 CH 4 is nonpolar: dispersion forces. S SO 2 SO 2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO 2 molecules.

56 Metallic METALLIC BONDING Bonding found in metals; holds metals together strongly. Electrons are free to move from atom to atom. Atoms are tightly packed together. Metals are: flexible, can bend and stretch, and can conduct electricity. Good conductors at all states, lustrous, very high melting points. Examples; Na, Fe, Al, Au, Co

57 METALLIC BONDING Electron Sea Model Explained by the Electron Sea Model The atoms in a metallic solid contribute their valence electrons to form a sea of electrons that surrounds metallic cations. Delocalized electrons are not held by any specific atom and can move easily throughout the solid. The more delocalized electrons the stronger the bond. A metallic bond is the attraction between these electrons and the metallic cation.

58 METALLIC BONDS Mellow dogs with plenty of bones to go around.

59 METALS FORM ALLOYS Metals do not combine with metals. They form alloys which is a solution of a metal in a metal. Examples are steel, brass, bronze and pewter.

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62 IONIC COMPOUNDS -Formed from a combination of metals and nonmetals. -Electron transfer from the cation to the anion. -Opposite charged ions attract each other. COVALENT COMPOUNDS -Formed from a combination of nonmetals. -Electron sharing between atoms. METALLIC COMPOUNDS -Formed from a combination of metals - sea of electrons ; electrons can move among atoms Solids at room temperature Can be solid, liquid, or gas at room temperature. Solids at room temperature High melting points Low melting points Various melting points Dissolve well in water Conduct electricity only when dissolved in water; electrolytes Do not dissolve in water (Sugar is an exception) Do not conduct electricity; non electrolytes Do not dissolve in water. Conduct electricity in solid form. Brittle, hard Soft Metallic compounds range in hardness. Group 1 and 2 metals are soft; transition metals are hard. Metals are malleable, ductile, and have luster.

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64 KEY IMPLICATIONS After bonding transpires, molecules and ions are formed. The resulting structure is a description of the packing and arrangements of these particle to give a substance (S/L/G). If the substance is a solid and the particles pack together in a regular repeating pattern then a crystal is formed. Both bonding and structure determines the physical properties of a substance.

65 LEARNING CHECK 1. Explain why ionic compounds do not conduct electricity in their crystalline form. 2. Why do metals and nonmetals usually form ionic compounds, whereas two bonded nonmetals are never ionic? Explain. 3. Why is the formation of ionic compounds exothermic? Describe whether the following compounds are likely to be ionic or not ionic based on the properties given. Explain your reasoning. a. Compound 1 has a melting point of 545 o C and dissolves in water. b. Compound 2 is a brittle material that is used to melt road ice during storms. 4. Why do ionic compounds tend to be hard? c. Compound 3 melts at 85 o C and catches fire when heated to 570 o C.