Intro To Chemistry Chpt 4 Chemical Bonds

Bonding pair:

The shared pair of electrons in the chlorine molecule is another example of a covalent bond; they are called a " ". Example in picture on the right

Polar:

δ⁺ H-Cl δ⁻ or +---> H--Cl

HONC Rule:

* Hydrogen forms one bond * Oxygen forms 2 bonds * Nitrogen forms 3 bonds * Carbon forms 4 bonds

Polar covalent bond:

An unequal sharing of of an electron pair (or pairs) leads to bond polarity. " " is not an ionic bond. In a " " , the atom at the positive end of the bond (hydrogen in HCl) still has a partial share of the bonding pair of electrons.

Group 5A:

Tend to pick up 3 electrons to form 3- ions

Covalent bond:

The bond formed when atoms share electrons. (shared pair of electrons) A " " is usually represented as a dash, so that the Lewis dot structure for H₂ Becomes H-H

Core electrons:

Those in all the other shells are lumped together

Group 7A:

non metal atoms take one electron to form 1- ions

Molecular shape Linear:

# of bonded atoms=2 # of lone pairs=0 # of electron sets=2 Examples of molecules: BeCl₂ HgCl₂ CO₂ HCN

Molecular shape (2) Bent:

# of bonded atoms=2 # of lone pairs=1 # of electron sets=3 Examples of molecules: SO₂ O₃ Example of SO₂ in picture on right

Electronegativity Differences: < 0.5 0.5-2.0 > 2.0

Bonding Types: Nonpolar covalent Polar covalent Ionic

Names of covalent compounds:

Covalent, or molecular, compounds are those in which electrons are shared, not transferred Such compounds generally have molecules that consist of 2 or more nonmetal atoms. Many covalent compounds have common and widely used names. Examples are: Water (H₂O) Ammonia (NH₃) Methane (CH₄) However, when naming most covalent compounds, systematic rules apply. The prefixes mono-, di-, tri, and so on are used to indicate the number of atoms of each element in the molecule. For example: The compound N₂O₄ is called dinitrogen tetroxide (the ending vowel is often dropped from tetra- and other prefixes when they precede another vowel.) We leave off the prefix mono- in front of the first atom in a formula but include it in front of the second atom (for instance, NO is nitrogen monoxide).

Valence electrons:

Electrons in the outermost shell The number of valence electrons of an element can be determined by the periodic table group of the element.

Group 2A:

In following the octet rule atoms of group " " give up 2 electrons to form 2+ ions

Group 3A:

In following the octet rule atoms of group " " give up 3 electrons to form 3+ ions

Group 4A:

In following the octet rule atoms of group " " give up 4 electrons to form 4+ ions

Group 1A:

In following the octet rule atoms of group " " give up one electron to form 1+ ions

Duet rule:

In the case of helium, a maximum of 2 electrons can occupy its single electron shell. Consequently, such small atoms as hydrogen (in metal hydrides) or lithium, often follow the " rule of two" or the "duet rule" To obtain the electron configuration of their nearest noble gas-(helium)

> 2.0

Ionic bond

H-H and Cl-Cl

Nonpolar

< 0.5

Nonpolar covalent bond

Ionic bond:

One atom completely loses an electron

Example of Lewis dot symbols for a compound:

See the picture to the right

Binary:

Simple ions of opposite charge can be combined to form " " (2 element) ionic compounds

Group 6A:

Tend to pick up 2 electrons to form 2- ions

Double bond:

The covalent bond that forms from 2 pairs of electrons being shared between 2 atoms and is indicated by a double dash between atoms But can be shown in a Lewis dot structure as well. Example on the right Double bond Oxygen

The Octet Rule:

The drive for 8 Atoms are driven toward becoming ions to have 8 -an octet electrons in their valence shells. An octet of electrons is characteristic of noble gases except helium. When atoms react with each other,they often tend to attain this very stable electron configuration of a noble gas closest to them. Thus they are said to follow the octet, or the " rule of eight"

Lone pairs or nonbonding pairs:

The electrons that stay on one atom and are not shared Example in picture on the right

VSEPR Theory:

Valence Shell Electron Pair Repulsion Theory We can use Lewis formulas as the initial step and then apply (VSEPR) Theory to predict the arrangement of atoms about a central atom or a bond-single,double, or triple-between the central atom and another atom. The basis of the VSEPR theory is that electron sets arrange themselves as far away from each other as possible so as to minimize repulsions between these like-charged particles If you google this you will find it online, it's way too much to try and type

Single bond:

When one pair of electrons is shared

Molecular shape (1 ) Bent:

# of bonded atoms=2 # of lone pairs=2 # of electron sets=4 Examples of molecules: H₂O H₂S SCl₂

Molecular shape trigonal planar:

# of bonded atoms=3 # of lone pairs=0 # of electron sets=3 Examples of molecules: BF₃ AlBr₃ CH₂O

0.5--2.0

Polar covalent bond

Molecular shape Trigonal pyramidal:

# of bonded atoms=3 # of lone pairs=1 # of electron sets=4 Examples of molecules: NH₃ PCl₃

Molecular shape Tetrahedral:

# of bonded atoms=4 # of lone pairs=0 # of electron sets=4 Example of molecules: CH₄ CBr₄ SiCl₄

The art of deduction works something like this:

* Fact: Noble gases, such as helium, neon, and argon, are inert; they undergo few, if any, chemical reactions * Theory: The inertness of noble gases result from their electron configurations Each (except helium) has an octet of electrons in its outermost shell. * Deductions: Other elements that can alter their electron configurations to become like those of noble gases would become less reactive by doing so. For example: Sodium has 11 electrons, 1 of which is in the 3rd shell. Recall that electrons in the outermost shell are called valence electrons, while those in all the other shells are clumped together as core electrons The outermost shell is filled when it contains 8 electrons, 2 in the s orbital/6 in a p orbital The first shell is an exception: It only holds 2 electrons in the 1s orbital If the sodium atom got rid of its valence electron, its remaining core electrons would have the same electron configuration as an atom of the noble gas neon Similarly, if a chlorine atom gained an electron, it would have the same electron configuration as argon. The drive to have an electron configuration like a noble gas is strong The sodium atom, having lost an electron becomes positively charged. It has 11 protons (11+) and only 10 electrons (10-). It is symbolized by Na⁺ and is called a sodium ion. The chlorine atom, having gained an electron, becomes negatively charged. It has 17 protons ( 17+) and 18 electrons (18-) It is symbolized by Cl⁻ and is called chloride ion *Note* that a positive charge, as in Na⁺, indicates that one electron has been lost. Similarly, a negative charge, as in Cl⁻ and Ar are isoelectronic (i.e., they have the same electron configuration), they are different chemical species, because their nuclei are different. In the same way, a sodium atom does not become an atom of neon when it loses an electron; the sodium ion simply is isoelectronic with neon.

Triple bond:

A covalent bond that forms from 2 atoms that share 3 pair of electrons and is indicated by a triple dash between atoms But can be shown in a Lewis dot structure as well. Example on the right Triple bond Nitrogen

Dipole:

A molecule with a positive end and a negative end, the polarity is commonly indicated with an arrow with a plus sign at the tail end The plus sign indicates the part of the molecule with a partial positive charge, & The head of the arrow signifies the end of the molecule with a partial negative charge.

Lewis (Electron-Dot) Symbols:

A way of representing atoms or molecules by showing electrons as dots surrounding the element symbol. One bond is represented as two electrons. For example: H. Other examples of Lewis dot symbols for other elements in the picture to the right.

Electronegativity:

An elements can be viewed as its ability to attract electrons in a molecule. The more electronegative an atom is, the greater its tendency to pull the electrons toward its end of the covalent bond. The atoms to the right in the periodic table are, in general, more electronegative than those to the left. The atoms on the on the right are precisely the atoms that, in forming ions, tend to gain electrons/form negatively charged ions (anions). The atoms on the left-metals-tend to give up electrons/become positively charged ion (cations) Within a column, electronegativity tends to be higher at the top and lower at the bottom of the column.

Polyatomic ions:

Are charged particles containing 2 or more covalently bonded atoms Many compounds contain both ionic and covalent bonds. There are many groups of atoms that (like hydroxide ion) remain together as ions, through most chemical reactions. Sodium hydroxide, commonly known as caustic soda or lye, consist of sodium ions (Na⁺) and hydroxide (OH⁻) The hydroxide ion contains an oxygen atom covalently bonded to a hydrogen atom, plus an "extra" electron, which gives it the negative charge. That extra electron is needed so that both the O and the H atoms have their valence shells filled. The chemical formula for sodium hydroxide is NaOH. For each sodium ion, there is one hydrogen ion, so that NaOH is neutral. The bond between the sodium and hydroxide ions is ionic, just like the ionic bonds you have already seen between sodium/chloride ions Example of sodium hydroxide using Lewis dot symbols in picture on the right

Polyatomic molecules:

Are electrically neutral groups of three or more atoms held together by covalent bonds. Molecules are distinguished from ions by their lack of electrical charge. The molecular formula for water is H₂O Oxygen needs 2 electrons to complete its octet, but each hydrogen only has one electron to share. Therefore, for oxygen to satisfy its octet, it must share electrons with 2 hydrogen atoms. This means that oxygen fulfills its octet by sharing 2 pairs of electrons One pair of electrons per hydrogen. This is how a polyatomic molecule is formed- a molecule containing more than 2 atoms Example of a water ( polyatomic) molecule formula in picture on the right

Group B metals (transitional metals):

Atoms of group " " can give up various numbers of electrons to form positive ions with various charges. (Many of the " " can form different ions with different charges)

Nonpolar covalent bonds:

Bonds where the electron pairs are shared equally; such as those you have already seen in the diatomic molecules of hydrogen, nitrogen, oxygen and chlorine.

Rules for writing Lewis formulas:

How to write " " for molecules or polyatomic ions. To write " " * First put the atoms of the molecule or ion in their proper places * Then place all the valence electrons so that each atom has a filled shell * The skeletal structure of a molecule tells us the order in which the atoms are attached to one another * Hydrogen atoms form only single bonds. They are always at the end of a sequence of atoms. Hydrogen is often bonded to carbon, nitrogen, or oxygen. * Oxygen tends to form 2 bonds * Nitrogen usually forms 3 bonds * Carbon forms 4 bonds (recall the HONC rule) * Polyatomic molecules/ions often consist of a central atom surrounded by atoms of higher electronegativity. (hydrogen is an exception;it is always on the outside, even when bonded to a more electronegative element.) * The central atom of a polyatomic molecule or ion is often the least electronegative atom. 1) Determine the total # of valence electrons. This total is the sum of the valence electrons for all the atoms in the molecule or ion. You must also account for the charge(s) on a polyatomic ion. a) For a polyatomic anion, add to its total # of valence electrons the # of negative charges. b) For a polyatomic cation, subtract the # of positive charges from the total # of valence electrons. Example: N₂O₄ has [2x5] + [4x6]=34 valence electrons NO₃⁻ has [1x5]+[3x6]+1=24 valence electrons NH₄⁺ has [1x5]+[4x1]-1=8 valence electrons 2) Write a reasonable skeletal structure & connect bonded pairs of atoms by dash (1 dash per shared electron pair) 3) Starting with the most electronegative atom in the polyatomic molecule or ion, place electron in pairs around outer atoms so that each has an octet, except for hydrogen, which will have a duet. 4) Subtract the # of electrons assigned so far (both in bonds/as lone pairs) from the total calculated in step 1. Any electrons that remain are assigned in pairs to the central atom(s) 5) If a central atom has fewer than 8 electrons after step 4, one or more multiple bonds are likely. Move one or more lone pairs from an outer atom to the space between the atoms to form a double or triple bond. A deficiency of 2 electrons suggests a double bond, and a shortage of 4 electrons indicates either 2 double bonds or a triple bond to the central atom. Example: Carbon dioxide (CO₂) 1). There are 4+(2x6)=16 valence electrons 2). The skeletal structure is O-C-O 3). First, we place 3 lone pairs on each oxygen atom O-C-O just imagine there are 3 lone pairs of electrons on each oxygen atom (i'm not able to type them there) 4). We have assigned 16 electrons. None remains to be placed. 5) The central carbon atoms has only 4 electrons. It needs to have 2 double bonds to achieve an octet. Move a lone pair from each oxygen atom to the space between the atoms to form a double bond on each side of the carbon atom Look at the picture on the right to see what it is supposed to look like in the end

Crossover method:

Write the formula for calcium chloride Solutions: first, write the symbols for the ions write the charge on chloride ion explicitly as "1-" to illustrate the method. You may omit the "1" when your are comfortable with the process Ca²⁺ Cl¹⁻ Cross over the charge numbers (without the charges) as subscripts Then write the formula. The formula for calcium chloride is Ca₁Cl₂ or (dropping the "1") simply CaCl₂