Atoms and BondingCh 2 and 3 Bettelheim text

Atoms and Bonding-Ch. 2 and 3 Bettelheim text **Recommended sites: kentchemistry. com Also, Janet Coonce videos on youtube • Atom-composed of p+, n 0 and e • Know atomic theory!

• • Diameter of a nucleus is only about 10 -15 m. Diameter of an atom is only about 10 -10 m. Fig 3. 1 The structure of an atom General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

Atomic number (Z) • • Always equals the p+ number In a neutral atom, p+=e. Mass number: p+ and n 0 Isotopes-different number of neutrons, so different masses

Isotopes of hydrogen General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

Periodic Table • Mendeleev-1869 -organized elements by atomic mass; forerunner of modern PT • Periods (rows) • Groups (columns) • Know names of groups • Metals, metalloids and non metals

• Fig 3. 2 The periodic table General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

Periodicity • Atomic radius-decreases across periods and increases down groups • Know characteristics of groups • Properties of elements determined by earrangement • Electronic structure (electron configurations) • *note-shell=principle energy level • S=spherical orbitals; P=dumbbell shaped

Electronic structure con’t • The maximum # of e- that can occupy an energy level (shell) is 2 n 2

• • The first shell has only an s subshell The second shell has and p subshell The third shell has an s, p, and d subshell. The fourth shell has an s, p, d, and f subshell. General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

• Electronic configuration of a few elements are shown below: General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

Sublevel Diagram, Order in which Sublevels are Usually Filled 1 s 2 s 2 p 3 s 3 p 3 d 4 s 4 p 4 d 4 f 5 s 5 p 5 d 5 f 6 s 6 p 6 d 7 s 7 p

Electrons in Energy Levels, Sublevels, and Orbitals Principal Available Energy Sublevels Level (n) (s, p, d, f) Orbitals in Sublevel (boxes) Possible Electrons in Sublevel Possible Electrons in Energy Level 1 s 1 2 2 2 s p 1 3 2 6 8 3 s p d 1 3 5 2 6 10 18 4 s p d f 1 3 5 7 2 6 10 14 32

Electron configurations con’t • In class- do configurations for C, N and P • Also, draw the electron config for potassium and the potassium ion • How might an abbreviated configuration of potassium be written?

General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 3

Valence Electrons - are the electrons in the outermost orbital of an atom - are associated with atom’s highest principal energy level (period) - only valence electrons determine chemical properties of an element.

Ionic Bonds • Typically, atoms from groups IA and IIA ionically bond with VIA an VIIA • Ion- Charged atom

The symbol for cation is written by adding a positive charge as a superscript to the symbol for the element. For example, by losing an electron Na metal is converted to the sodium cation (Na+). General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 4

• Gain of one or more electrons by a neutral atom gives a negatively charged ion called a anion. The symbol for the anion is written as by adding a negative charge as a superscript to the symbol for the element. For example, by gaining an electron chlorine atom becomes chloride ion, Cl-. General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 4

• Halogens: Large ionization energy – electron not easily lost; Large electron affinity – electron easily gained – formation of anion is favored. • Fig 4. 1 Relative ionization energies (red) and electron affinities (blue) for elements in the first four rows of the periodic table. General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 4

Where do the electrons come from? Where do the electrons go? • Atoms with unfilled valence shells can become filled by transferring or sharing electrons with other atoms • When atoms come together by sharing or transferring electrons, a chemical bond results. • After a chemical bond is formed, the atoms involved in the bond each have a filled valence shell (octet).

• • In Na. Cl crystal, each sodium ion is surrounded by 6 chloride ions, and each chloride ion is surrounded by 6 sodium ions. In a Na. Cl solution, individual sodium and chloride ions are surrounded by water. +H 2 O Notice the orientation of the water molecules

Ionic Bonds • • Opposite electrical charges attract each other. When sodium combines with chlorine, sodium transfer electron to chlorine forming Na+ and Clions. The oppositely charged Na+ and Cl- ions are held together by a ionic bond. In three dimensional crystals, all nearby oppositely charged ions attract each other. We can not separate attraction between only two oppositely charged ions -we speak of the whole crystal and crystal as being an ionic solid or ionic compound. (molecule)

Some Properties of Ionic Compounds • Ionic compounds are usually crystalline solids. • Different ions vary in size and charge. • Ions packed together in crystals in different ways. • The crystal pattern ensures that ions efficiently fills space and maximizes ionic bonding. • Ionic compounds have high melting and high boiling points.

• The formula of an ionic compound shows the lowest possible ratio of atoms in the compound.

Naming ionic compounds • We will NOT cover the “old” naming system • For binary ionic compounds, name the cation first and then the anion. However, you must change the ending of the anion to -ide • If a transition metal is the cation, a roman numeral must be given to indicate it’s oxidation number (charge) • Although polyatomic ions bond covalently, they may ionically bond with a cation or an anion • For test, know biological importance of Ca, Fe, K and Na • Polyatomic ions to know for this test: See chart given in class and know the ones I told you to check [especially acetate, hydroxide, nitrate, carbonate, sulfate, phosphate, ammonium, cyanide, chromate, dichromate, hydrogen carbonate (bicarbonate)]. Quiz Monday!

Practice- Write formulas or name: • • 1. Sodium carbonate 2. Calcium phosphate 3. K 2 SO 4 4. Fe. Cr 2 O 7

Acids and bases • One definition (Arrhenius): Acid-substance that donates H+ in water (ex: HCl) Base- donates OH- (ex: Na. OH) **Bronsted-Lowry definition: Acid is p+ donor; base is p+ acceptor (ex: NH 3 + H 2 O ? ) • Hydronium ion H 30+ • Exp: HCl, Na. OH • p. H scale 0 -7 -14 • p. H + p. OH = 14; p. H= -log[H+] • Properties of acids and bases (Acids- react with metals; “sour”; Bases-slippery, bitter) • Acid + Hydroxide Base Salt and Water

Acids and bases • Problem with Arrhenius definition: Not all acid/ base reactions occur in water • Bronsted- Lowry definition: Acid is a proton (H+) donor and base is proton acceptor • NH 3 + HCl-> NH 4+ + Cl-

More acids and bases • Buffer-substance that can accept hydrogen or hydroxide ions in solution. Prevent drastic changes in p. H • Major buffers in the body: 1) bicarbonate (hydrogen carbonate); 2)hydrogen phosphate • CO 2 + H 2 O H 2 CO 3 H+ + HCO 3 -

Covalent Bonds • • • Covalent bond: A bonds formed by sharing electrons between atoms. Usually nonmetals Molecule: A group of atoms held together by covalent bonds. The nonmetals near the middle of the periodic table reach an electron octet by sharing an appropriate number of electrons. General, Organic, and Biological Chemistry; Prentice Hall @ 2003, Ch. 5

• Spherical 1 S orbital of two individual hydrogen atoms bends together and overlap to give an egg shaped region in the hydrogen molecule. The shared pair of electrons in a covalent bond is often represented as a line between atoms. *Remember, hydrogen doesn’t get a full octet, it only fills the first shell s orbital

• A water molecule results when two hydrogen atoms and one oxygen atom are covalently bonded in a way shown in the following picture: How many electrons are in the valence shell of each atom in water?

Let’s go back to the electron configurations: Hydrogen atom: 1 proton, 1 electron. 1 s 1 Oxygen atom: 8 protons, 8 neutrons, 8 electrons. 1 s 2 2 p 4 Which electronic orbitals are involved in the bonds? Hydrogen: 1 s Oxygen: 2 p

Fig 5. 4 For P, S, Cl, and other elements in the third period and below, the number of covalent bonds may vary, as indicated by the numbers shown in parentheses.

Numbers of Covalent Bonds - single bond– sharing of two electrons - longest bond length - double bond- sharing of four electrons - middle bond length - triple bond- sharing of six electrons - shortest bond length - strongest covalent bond - The electrons in the valence shell that are not shared are called lone pairs.

Drawing Lewis Structures • To draw Lewis structure, you need to know the connections among atoms. Also, common bonding patterns, shown below, for C, N, O, X, and H simplifies writing Lewis structure.

Method #1 - See next slides and/or handout. I PREFER THIS WAY Method #2 - Book’s suggested rules for drawing molecules (Lewis structures)**note-these don’t always work…they are merely a guide. Some have expanded octets. 1. Determine the number of valence electrons in the molecule. 2. Determine the arrangement of atoms in the molecule (least electronegative –center) 3. Connect the atoms with single bonds 4. Give each peripheral atom an octet by adding lone pairs 5. Place all remaining electrons on central atom by adding lone pairs 6. If the central atom still doesn’t have an octet, take a lone pair from a peripheral atom to form a multiple bond

LEWIS DOT STRUCTURES 1. Arrange the symbols such that the least electronegative element is in the center and the other elements are surrounding the central atom. O C O 2. Give each of the elements their appropriate number of valence electrons (dots). Remember the number of valence electrons for a representative element is the same as the group number. . . : O : C: O: . . 3. Keep track of the total numbr of valence electrons for the compound by adding the valence electrons from each atom. If the compound is an ion then add electrons (dots) for each negative charge or subtract electrons (dots) for each positive charge. 4 for C and 6 for O (twice) = 16 electrons

LEWIS DOT STRUCTURES 4. Now move the dots around so that you have 8 dots (the octet rule) around each element (do not forget the exceptions) while at the same time keeping the dots in pairs. Electrons, at this point, exist as pairs (the buddy system). 5. EXCEPTIONS TO THE OCTET RULE: Group I, II, and III need only 2, 4, and 6 electrons, respectively, around that atom (if they are covalently bonding). 6. If there are too few pairs to give each atom eight electrons, change the single bonds between two atoms to either double or triple bonds by moving the unbonded pairs of electrons next to a bonding pair. . . : O: : C: : O:

LEWIS DOT STRUCTURES 7. Once the octet rule has been satisfied for each atom in the molecule then you may replace each pair of dots between two atoms with a dash. . . : O =C= O: 8. Now check your structure by a) count the total number of electrons to make sure you did not lose or gain electrons during the process. b) Use FORMAL CHARGE (FC) calculations as a guideline to the correct structure. A zero formal charge is usually a good indication of a stable structure. FC (X) = # of valence electrons - (1/2 bonding electrons + nonbonding electrons) For our example: FC(C) = 4 - (1/2 8 + 0) = 0 FC(O) = 6 - (1/2 4 + 4) = 0

The Basics: Drawing Lewis Structures Step 1: Calculate the total number of valence electrons in the molecule or ion Step 2: Determine the central atom(s) of the molecule or ion – usually it’s the least electronegative atom. Step 3: Draw a tentative diagram for the molecule or ion. Rules a) A hydrogen atom always forms one bond. Hydrogen is always a terminal atom in a Lewis diagram – an atom that is bonded to only one other atom. b) A carbon atom normally forms four bonds c) When several carbon atoms appear in the same molecule, they are often bonded to each other.

Draw the Lewis Structure for the following molecules. 1. H 2 O 2. CO

Draw the Lewis Structure for the following molecules. 3. BH 3 4. NH 3

CW- Draw CCl 4

CW- Draw PBr 3

Draw OCl 2

Draw NH 4+ *Cation **Coordinate covalent bond- one atom donates both shared e-

Explain the last example in terms of acid/base • NH 3 is the ? • H+ is the ?

Classwork: • Do “WS Covalent Lewis Dot Structures”

The next few slides have…. • Some more “challenging” Lewis structures. I will let you know if we are doing them or not.

Draw SO 3 2*anion; name it!

Draw SO 2 *Resonant structure. Draw both

Practice: • Predict the most stable structure: ONC- or OCN- or NOC-

LEWIS DOT STRUCTURES Predict the most stable structure: ONC- or OCN- or NOC. . . : O: : N: : C: or : O: : C: : N: or : N: : O: : C: 1) Total electrons is: 6 e- for O + 5 e- for N + 4 e- for C + 1 e- for negative charge = 16 etotal. All structures fulfill the octet rule. 2) FC (X) = # of valence electrons - (1/2 bonding electrons + nonbonding electrons) structure#1: FC(C) = 4 - (1/2 4 + 4) = -2 FC(O) = 6 - (1/2 4 + 4) = 0 FC(N) = 5 -(1/2 8 + 0) = +1 structure#2: FC(C) = 4 - (1/2 8 + 0) = 0 FC(O) = 6 - (1/2 4 + 4) = 0 FC(N) = 5 -(1/2 4 + 4) = -1 structure #3: FC(C) = 4 - (1/2 4 + 4) = -2 FC(O) = 6 - (1/2 8 + 0) = +2 FC(N) = 5 -(1/2 4 + 4) = -1 structure #2 has the combination with the lowest formal charge. It also has the negative formal charge on one of the more electronegative atoms Calculate the formal charge for the most stable structure: . . : O: C: : : N: . . (-1, 0, 0)

Practice problems • Draw: Cl. O 2 • Si. H 4 • As. H 3

More Practice Problems Cl. O 2 - As. H 3 Si. H 4

1. Group Study Problems Draw the Lewis structure for the following a) H 2 S b) PH 3 c) CH 2 O d) NO 2 - e) H 2 CO 3 f) CBr 4 g) CH 2 FCl h) C 2 H 2 I) O 3 2. Calculate the formal charge for “c”, “d”, “f”, and “I”.

So, how do you draw these structures? * Practice any we didn’t do in class at home • Br 2 • CH 4 • HCl • Si. F 4 • PF 3 • H 2 CO • C 2 H 6 • C 2 H 4 • CO 2 • CO • H 2 O • OH • NH 3 • HCN • NH 4 + (What is a coordinate covalent bond? • NO 3 - (*use the ‘rules” for this one. Also, this one displays resonance. What is that? ) • CO 3 2 • SO 2

Shape of Molecules • Molecular shapes can be predicted by noting how many bonds and electron pairs surround individual atoms and applying what is called the valence-shell electron-pair repulsion (VSEPR) model. The basic idea of VSEPR model is that the negatively charged clouds of electron in bonds and lone pair repel each otherefore tends to keep apart as far as possible causing molecules to assume specific shape.

• • There are three step to applying the VSEPR model: Step 1: Draw molecule, and identify the atom whose geometry is of interest. Step 2: Count number electron the ofcharge clouds surrounding the atom of interest. Step 3: Predict molecular shape by assuming that the charge clouds orient in space so that they are as far away from one another as possible.

VSEPR Model The shape depends on the number of charged clouds surrounding the atom © 1995 -2002 by Prentice-Hall, Inc. A Pearson Company

Bond angles to know: • Linear-180 • Trigonal planar-120 (If all e- are in bonded pairs like BF 3. If one pair is lone…. and thus “bent”, like SO 2 , then it’s actually 119) • Tetrahedral-109. 5 • Pyramidal-107 (lone pair pushes bonded pairs closer together) • Bent (water)-104. 5 (2 lone pairs) • Practice: http: //www. chem. umass. edu/genchem/wh elan/Class_Handouts/111_Molecular_Geometry_ Worksheet. pdf

Polar Covalent Bonds and Electronegativity • • • Electrons in a covalent bond occupy the region between the bonded atoms. If the atoms are identical, as in H 2 and Cl 2, electrons are attracted equally to both atoms and are shared equally. If the atoms are not identical, however, as in HCl, the bonding electrons may be attracted more strongly by one atom than by the other and may thus are shared unequally. Such bonds are known as polar covalent bonds.

• In HCl, electron spend more time near the chlorine than the hydrogen. Although the molecule is overall neutral, the chlorine is more negative than the hydrogen, resulting in partial charges on the atoms. • Partial charges are represented by placing d- on the more negative atom and d+ on the more positive atom. • Ability of an atom to attract electrons is called the atom’s electronegativity. • Fluorine, the most electronegative element, assigned a value of 4, and less electronegative atoms assigned lower values • Electroneg differences >1. 7 -1. 9 considered ionic; <1. 7 -1. 9 covalent

• Fig 5. 7 Electronegativities and the periodic table

Polar Molecules • • Entire molecule can be polar if electrons are attracted more strongly to one part of the molecule than to another. Molecule’s polarity is due to the sum of all individual bond polarities (sometimes called dipolar forces or dipoles) and lone -pair contribution in the molecule.

• Molecular polarity is represented by an arrow pointing at the negative end and is crossed at the positive end to resemble a positive sign.

• Molecular polarity depends on the shape of the molecule as well as the presence of polar covalent bonds and lone-pairs.

Naming Binary Molecular Compounds • • When two different elements combines together they form binary compound. The formulas of binary compounds are usually written with the less electronegative element first. Thus, metals are always written before non-metals. Prefix such as mono, di, tri, tetra etc, are used to indicate number of atoms of each element.

A few examples of binary compounds are given below:

• the • The following two steps guide is helpful in naming binary compounds: Step 1: Name formula, using a prefix if needed to indicate the number of atoms. ment second • Step the 2: Name formula, using an –ide ending as for anions, along with a prefix if needed to indicate the number of atoms.

Characteristic of Molecular Compounds • Molecules are neutral as a result there is no strong electrical attractions between the molecules to hold them together. However, there are several weaker forces exist between molecules, known as intermolecular forces (van der Waals forces). When • intermolecular forces very weak, are molecules are weakly attracted to one another and that the substance is gas at ordinary temperature.

• If the intermolecular forces are somewhat stronger, the molecules are pulled together into a liquid. • If the forces are stronger, the substance becomes a molecular solid. • Hydrogen bond-specific type of intermolecular force involving hydrogen (that is cov. bonded to a highly electronegative atom) attracted to an unshared electron pair of another molecule. Ex-water • solids are lower than those of ionic solids. • Most molecular compounds are insoluble in water. ounds ar • melted because they have no charged particles.

Remember • Electronegativity increases across periods and decreases down groups • Ionization Energy increases across periods and decreases down groups • Atomic radius decreases across periods and increases down groups

Linear Bent Trigonal planar Tetrahedral