CHEM 102 (Pre-Exam 1) by Christelle Seri

Groups of the Periodic Table I

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Stability of an Atom

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Groups of the Periodic Table II

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Let's play a game or just mess around if you want

Moles?

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K here goes
Avogadro's Number:
Multiply a given number of atoms (photons in certain cases) by this number to get moles.

Visible Light Spectrum

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What does it mean to say that the energy of the electrons in an atom is quantized?

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Electrons exist on specific energy levels and cannot exist in between them. When an electron moves to another energy level, it must release or absorb very specific amounts of energy. This is what quantized means: the amount of energy absorbed or released can only be specific quantities called quanta. @

Orbitals and the Periodic Table

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Watch out for the positioning of the d block

Orbital Shape

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Valence Electrons

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Group of the Periodic table is the same as the number of valence shell electrons (!!!) {Almost failed a problem on this}

Polyatomic ions (pt. 1

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Electron Affinity Values

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Negative values of EA means energy is released. Positive values of EA means energy is required for the electron to become negatively charged.

Electron Affinity and the Periodic Table

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As we go from left to right across a period EAs tend to become more negative.

Effective Nuclear Charge and its Implications

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As the effective nuclear charge increases, it becomes easier to add an electron.

As electrons are removed, the remaining core electrons experience a greater effective nuclear charge and are drawn closer to the nucleus.

Electron Affinity Exceptions

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group 2 (2A): have a completely filled ns subshell, electron must be added to a higher np subshell which is more difficult

group 15 (5A): have a half filled np subshell, electron must be paired with an existing np

group 18 (8A): have a completely filled shell, electron must be added to a higher n which is more difficult

Initial relative stability of the electron configuration disrupts the trend

group 17 (7A): have the most negative electron affinities

Electron Affinity and Groups of the Periodic Table

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Second element most often has the greatest EA.

Why?
Because the n=2 shell is relatively small resulting in large electron-electron repulsions.

Metallic Properties

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Metallic properties depend on having easily removable electron. So these increase down a group and decreases over the period.

Periodic Table Ion Predictions

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main-group metals: lose enough electrons to equal number of electrons in an atom of the preceding noble gas.

nonmetal ions: gain enough electrons to equal number of

electrons in an atom of the succeeding noble gas.

Moving from the far right to the left, elements often form anions with a negative charge equal to the number of groups moved left from the noble gases.

Moving from the far left to the right, main-group elements tend to form cations with a charge equal to the group number.

Predictive value decreases when moving toward the center of the periodic table

Periodic Table

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Polyatomic ions (pt. 2)

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Ionic Compounds

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When an element composed of atoms that readily lose electrons (a metal) reacts with an element composed of atoms

that readily gain electrons (a nonmetal), a transfer of electrons usually occurs, producing ions. The compound formed

by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the ions of opposite charge present

in the compound.

When a metal is combined with one or more nonmetals,

the compound is usually ionic.

Ionic Compound Properties

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solids that typically melt at high temperatures and boil at even higher temperatures
not electrically conductive in solid form because its ions are unable to flow
can conduct electricity when molten (or dissolved) because its ions are able to move freely through the liquid
exhibit a crystalline structure and tend to be
rigid and brittle
very strong however dissolve readily in water

Ionic Compound Ratios

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Ionic compounds are electrically neutral because total number of positive charges of the cations equals the total number of negative

charges of the anions.

The formula of an ionic
compound must have a ratio of ions such that the numbers of positive and negative charges are equal.

Molecular (Covalent) Compounds

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result when atoms share, rather than transfer (gain or lose), electrons
often exist as gases, low-boiling liquids, and low-melting solids under normal conditions
usually formed by a combination of nonmetals
poor conductors in any state

Atomic Radius

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The size of the atom (and its covalent radius) must increase as we increase distance.

Over a period, due to increasing effective charge the radius decreases.

Down a group, electrons are added to new shells which increases radius.

Ionization Energy

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The first ionization energy is the energy required to form a cation with +1 charge:

X(g)-> X⁺(g) + e^-

Values are always positive
As size increases the ionization energy decreases.
Increases over a period
Decreases down a group

Ionic Radius

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A cation has a smaller radius than the atom it was made from.

An anion has a larger radius than the atom it was made from.

Valence Electrons

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Valence electrons tend to have the highest energy and are the most likely to react.

Why is Electron Affinity Usually Negative?

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Cation Formation

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Most transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s electron(s) first
When forming a cation, an atom of a main group element tends to lose all of its valence electrons
The charge of a cation formed by the
loss of all valence electrons is equal to the group number minus 10

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Ag+, Zn2+, Cd2+

Order for Writing Elements ia Compound

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Diatomic Gases

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Naming

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For "Fun"

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MAHJONG!*

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*doesn't cover everything

Exothermic Electron Affinity

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Electronegatitivity

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Electronegativity generally increases going from left to right across a period and decreases going down a group for the representative elements. Use the difference in electronegativity to describe bond polarity. Don't forget that hydrogen has electronegativity equal to phosphorus and between boron and carbon.Remember, the greater the difference in electronegativity between the bonded atoms, the more polar the bond. Look at the Periodic Table.

Electronegativity Trends

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Consider the electronegativity trends.
The classifications ionic, polar covalent, and nonpolar covalent lie on a continuum.
In general, we can say that the bond is ionic if the difference in electronegativity is
large. This occurs when the bond involves a metal and a nonmetal. At the other end of the continuum is nonpolar covalent, when the difference in electronegativity
is essentially zero. This occurs when the bond involves two identical nonmetal atoms,
as well as bonds involving hydrogen bonded to carbon, boron, or phosphorus. All the
rest of the bonds between two nonmetals are assumed to lie between the two extremes
and can be classified as polar covalent. Look at the Periodic Table

Metals

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Metals lose electrons, to form cations.
Group I metals (Li → Fr, 1 valence electron) lose 1 electron.
This leaves them in the 1+ state.

Group II metals (Be → Ra, 2 valence electrons) lose 2 electrons.
This leaves them in the 2+ state.

Most of the Transition metals can exist in a number of different states.

For example, manganese, Mn, can exist in the 2+, 3+, 4+, 6+ and 7+ states.
To specify a state, it is written in Roman Numerals, directly after the metal name.
So, Mn⁴⁺ would be written as manganese(IV).

There are some Transition metals with more limited stable states.
Silver, Ag, is only stable as Ag⁺.
Zinc, Zn, is only stable as Zn²⁺.
Cadmium, Cd, is only stable as Cd²⁺.
No Roman Numerals are used in these cases, since there is only one possible state.

Aluminum, Al, is only stable as Al³⁺. The Roman Numeral is again, not needed.

Nonmetals

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When individual non-metallic elements react with metals to form salts, they gain electrons, forming anions.

Group VII elements (F → I, seven valence electrons) gain 1 electron to form the 1- state.
They are given the ending -ide, like fluoride (F^-), chloride (Cl^-), bromide (Br^-) and iodide (I^-).

Group VI elements (O → Se, six valence electrons) gain 2 electrons to form the 2- state.
They are again given the ending -ide, like oxide (O^2-).

Group V elements (N and P, five valence electrons) gain 3 electrons to form the 3- state and have the ending -ide (N^3- is nitride).

Carbon, with four valence electrons, can gain 4 electrons to form carbide, C^4-.

Hydrogen

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Hydrogen is a special case. It can gain an electron, to form H^-. This is called hydride and happens when hydrogen is reacting with a metal.
Hydrogen can also lose an electron to form H⁺, a proton, when reacting with non-metals.

Binary Salts

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Binary salts, which contain only two elements, are named for the metal and non-metal which combine to form them. So, when sodium and chlorine react, they form sodium chloride. The sodium atom loses an electron to form the Na⁺ ion, and chlorine gains an electron to form the Cl^- ion. Anions (negatively charged) and cations (positively charged) always combine to give neutral species, i.e., the number or positive and negative charges will be equal. So, the formula for sodium chloride is NaCl, with one positive and one negative charge. The formula of magnesium chloride is MgCl₂. Each chloride has a charge of 1- and the magnesium ion has a charge of 2+.
Write the formulas for the salts made by combining the following elements. Don't worry about subscripts, MgCl₂ would be written as MgCl2.

Naming Acids and Bases

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Isoelectronic Series

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An isoelectronic series is a group of atoms/ions that contain the same number of electrons.

HYDROGEN WTF

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Because hydrogen only needs

two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule.

Lewis Structure for Complex Molecules and Ions

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For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:

1. Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.

2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).

3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet

around each atom.

4. Place all remaining electrons on the central atom.

5. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain

octets wherever possible.

The Screw Ups

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Hydrogen is almost never a central atom in the Lewis Structure (an anomaly).
As the most electronegative element, fluorine also cannot be a central atom.
Boron tends to form compounds with fewer than eight electrons

Exceptions to the Octet Rule

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Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:

• Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.

• Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration. Generally,

these are molecules with central atoms from groups 2 and 12, outer atoms that are hydrogen, or other atoms that do not

form multiple bonds

• Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.

Hypervalent Molecules

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Elements in the third and higher

periods (n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called

hypervalent molecules.

Formal Charge

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The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms..

Thus, we calculate formal charge as follows:

formal charge = # of valence shell electrons - # of lone pairs - 1/2 # of bonding electrons

Double Checking Formal Charge

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We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.

Guidelines for Molecular Structure

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1. A molecular structure in which all formal charges are zero is preferable to one in which some formal charges

are not zero.

2. If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.

3. Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.

4. When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

A Comment on Double Bonds

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A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms.

Resonance

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If two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by

the various Lewis structures.

More on Resonance

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We call the individual Lewis structures resonance forms. The

actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the

individual resonance forms.

Supplemental Lewis Structure Resource

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Species with Expanded Shells

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These species have more than 8 electrons around the central atom. These structures are only possible when the principal quantum number is greater than n = 3 and d orbitals are available for bonding. These are most common with large central atoms bonded to small, electronegative atoms.

VSEPR Theory

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Electron Pair Repulsions

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From greatest to least:
1. lone pair-lone pair
>
2. lone pair-bonding pair
>
3. bonding pair-bonding pair

Order of Sizes

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From largest to smallest:
1. lone pair
>
2. triple bond
>
3. double bond
>
4. single bond

Electron Pair Geometries

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Using VSEPR Theory

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1. Write the Lewis structure of the molecule or polyatomic ion.

2. Count the number of regions of electron density (lone pairs and bonds) around the central atom. A single, double, or triple bond counts as one region of electron density.

3. Identify the electron-pair geometry based on the number of regions of electron density: linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral (Figure 4.19, first column).

4. Use the number of lone pairs to determine the molecular structure (Figure 4.19). If more than one

arrangement of lone pairs and chemical bonds is possible, choose the one that will minimize repulsions, remembering that lone pairs occupy more space than multiple bonds, which occupy more space than single bonds. In trigonal bipyramidal arrangements, repulsion is minimized when every lone pair is in an equatorial position. In an octahedral arrangement with two lone pairs, repulsion is minimized when the lone pairs are on opposite sides of the central atom.

Dipole Moment Formula

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μ= Qr

where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between

the charges

Diatomic Molecules

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For diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity

Polar Molecules

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To summarize, to be polar, a molecule must:

1. Contain at least one polar covalent bond.

2. Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.

Properties of Polar Molecules

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Polar molecules tend to align when placed in an electric field with the positive end of the molecule oriented toward the negative plate and the negative end toward the positive plate (Figure 4.28). We can use an electrically charged object to attract polar molecules, but nonpolar molecules are not attracted. Also, polar solvents are better at dissolving polar substances, and nonpolar solvents are better at dissolving nonpolar substances.

"Studying"

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"Studying" (Pt. 2)

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Covalent Bond Requirements

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an orbital on one atom overlaps an orbital on a second atom
the single electrons in each orbital combine to form an electron pair

Overlap

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orientation of orbitals and distance affects their overlap
energy of the system depends on how much the orbitals overlap
strength of a covalent bond depends on the extent of overlap of the orbitals involved

Sigma (σ) bonds

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Form from the overlap of:

two s orbitals
an s orbital and a p orbital
two p orbitals

Sigma and Pi Bonds

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While all single bonds are σ bonds, multiple bonds consist of both σ and π bonds.

Between any two atoms, the first bond formed will always be a σ bond, but there can only be one σ bond in any one location. In any multiple bond, there will be one σ bond, and the remaining one or two bonds will be π bonds.

Understanding Hybridization

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The following ideas are important in understanding hybridization:

Hybrid orbitals do not exist in isolated atoms. They are formed only in covalently bonded atoms.
Hybrid orbitals have shapes and orientations that are very different from those of the atomic orbitals in isolated atoms.
A set of hybrid orbitals is generated by combining atomic orbitals. The number of hybrid orbitals in a set is equal to the number of atomic orbitals that were combined to produce the set.
All orbitals in a set of hybrid orbitals are equivalent in shape and energy.
The type of hybrid orbitals formed in a bonded atom depends on its electron-pair geometry as predicted by the
VSEPR theory.
Hybrid orbitals overlap to form σ bonds. Unhybridized orbitals overlap to form π bonds.

Steps of Hybridization

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Determine the Lewis structure of the molecule.
Determine the number of regions of electron density around an atom using VSEPR theory, in which single bonds, multiple bonds, radicals, and lone pairs each count as one region.
Assign the set of hybridized orbitals from Figure 5.21 that corresponds to this geometry.

Figure 5.21

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christouseri — 2018-02-13 00:33:15 UTC

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Hybridization?

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Hydrogen....

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Hydrogen forms covalent bonds

DON'T FORGET

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Transition metals lose their s electrons before the d electrons

Noble Gases and Electronegativity

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Noble gases do not have electronegativity, they are an exception to the periodic trend

Nonmetal Properties

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High ionization energy and high electronegativity

Definition electronegativity v. electron affinity

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Fluorine has larger electronegativity than hydrogen

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Electronegativity is defined as the ability of an atom involved in a chemical bond to polarize electron density towards itself. Electronegativity tends to increase across a Period (from left to right as we face the Table) and decrease down a Group.

Incomplete electron shells tend to shield nuclear charge very imperfectly. This is manifested by the well-known contraction in atomic radii as we go from element to element across a Period from left to right.

And thus fluorine, to the right of the second Period, Z=9, is more electronegative than hydrogen, Z=1.

Shielding Effect and the Periodic Table

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Shielding Effect and Metals/Non-Metals

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