Learning Objectives Study Set- Chem Exam 2

Atomic radii decrease from left to right in a period (Na → Ar) on the periodic table. Choose the best explanation for this observed trend.

As the effective nuclear charge (Zeff) increases, all electrons in the outer shell (in this case the third shell) are attracted more strongly to the nucleus and are pulled closer to the nucleus.

Atom vs Element

Atoms make up elements Explanation: An atom is the part of an element. A particular element is composed of only one type of atom. Atoms are further composed of subatomic particles called electrons, protons and neutrons. Elements can combine with each other to form molecules via chemical reaction.

radii comparisons

Cations are always smaller than the neutral atom from which it is derived. Anions are always larger than the neutral atom from which it is derived. Therefore, the smaller ion in a pair will be the one with fewer electrons.

ionic bonding

Chemical bonding that results from the electrical attraction between cations and anions

chemical vs empirical formula

Chemical formulas tell you how many atoms of each element are in a compound, and empirical formulas tell you the simplest or most reduced ratio of elements in a compound. If a compound's chemical formula cannot be reduced any more, then the empirical formula is the same as the chemical formula (as is the case in ionic compounds).

nonpolar covalent bond

Electrons are shared equally

Coulomb's Law

F= kq_1q_2/r^2 k= constant q_1 and q_2= point charges r= distance between the point charges (change from pm to m)

Interruptions in Ionization Energy trends pt. 2

Google explanation: There are however some exceptions across every period where the ionization energy drops between an atom of group 2 and group 3 (like Mg and Al) and between group 5 and group 6 (like P and S). 1. drop in group 2 and group 3 (like Mg and Al) is due to a slight increase in distance from nucleus as outer electron occupies a new subshell (p subshell) slightly further away from the nucleus. 2. drop in group 5 and group 6 (like P and S) is due to spin pair repulsion which is due to the presence of 2 electrons in the same p orbital. This makes it require less energy to remove.

pure covalent bond

Neutral atoms held together by equally shared electrons

When talking about lattice energy....

Only use the term ions. not molecules or atoms

ionic bond

Oppositely charged ions held together by electrostatic attraction

Lone pairs

Pairs of valence electrons not involved in bonding; shown as pairs of dots on individual atoms

polar covalent bond

Partially charged atoms held together by unequally shared electrons

ground state of an atom

State of lowest energy, in which all electrons are in the lowest possible orbitals

ionic bonding is.....

The energy change for electron transfer results in a large gain in stability when the ionic bond forms. long explanation: electron transfer; refers to the electrostatic attraction that holds oppositely charged ions together in an ionic compound. *The attraction between the cation and anion draws them together

How does Coulomb's Law explain the force of attraction between the nucleus and any electron?

The force (F) in Coulomb's Law is analogous to energy (E) of an electron when that electron is interacting with a nucleus. As two particles are brought closer together, the value of F (force or energy) becomes more negative, which means that the particles are more attracted to each other.

Why does ionization energy increase across a period?

The ionization energy increases across the periods because the nuclear charge of the nucleus is increasing, and the atomic radius is decreasing. These two things increase the pull on the electrons from the nucleus, which makes it require more energy to remove an electron.

Charges in the periodic table

+1, +2, +3, +/-4, -3, -2, -1, 0 -Group 4-7, look at the elements at the top of the row -Groups 1-3, look at the elements right below the top of the row -Group 8 look at Ne (2nd to top)

How to distinguish between nonpolar covalent and polar covalent or between polar covalent and ionic

-A bond between atoms whose electronegativities differ by less than ~0.4 is generally considered purely covalent or nonpolar -A bond between atoms whose electronegativities differ by the range of ~0.4 to 2.0 is generally considered polar covalent -A bond between atoms whose electronegativities differ by 2.0 or more is generally considered ionic

Exceptions to the Octet Rule

-Incomplete octets: Some stable compounds have less than 8 electrons around the central atom (Elements in Groups 2A and 3A) -Odd numbers of electrons: Octet rule cannot be obeyed if there is an odd number of electrons (Free radicals or radicals) -Expanded octets: Elements in the third period and beyond can have more than 8 valence electrons because the 3d is now available.

two things that affect lattice energy

-charge -distance/size

Two extremes in the spectrum of bonding

-covalent bonds occur between atoms that share electron -ionic bonds occur between a metal and a nonmetal and involve ions *polar bonds fall between these extremes

electronegativity trend

-increases up and to the right -decreases down a group; increases from left to right in a period

For a given pair of atoms...

-triple bonds are shorter than double bonds -double bonds are shorter than single bonds

When there is more than one possible structure, the best arrangement is determined by the following guidelines

1) A Lewis structure in which all formal charges are zero is preferred 2) Small formal charges are preferred to large formal charges 3) Formal charges should be consistent with electronegativities.

To determine associated electrons

1) All the atom's nonbonding electrons are associated with the atom 2) Half the atom's bonding electrons are associated with the atom.

Steps to draw lewis structures

1) Draw the skeletal structure of the compound. The least electronegative atom is usually the central atom. Draw a single covalent bond between the central atom and each of the surrounding atoms. (H cannot be central!) 2) Count the total number of valence electrons present; add electrons for negative charges and subtract electrons for positive charges. 3) For each bond in the skeletal structure, subtract two electrons from the total valence electrons. 4) Use the remaining electrons to complete octets of the terminal atoms by placing pairs of electrons on each atom. Complete the octets of the most electronegative atom first. 5) Place any remaining electrons in pairs on the central atom. 6) If the central atom has fewer than eight electrons, move one or more pairs from the terminal atoms to form multiple bonds between the central atom and terminal atoms.

Interruptions in Ionization Trends

1. (between groups 2a AND 3a) electron removed from p orbital rather than s orbital. small amount of repulsion by s electrons 2. (between groups 5a and 6a) electron removed comes from doubly occupied orbital. repulsion from other electron in orbital helps in its removal instead of increasing from left to right, if an atom decreases left to right it's because an electron in a doubly occupied orbital is repelled by the other electron and requires less energy to remove than an electron in a half-filled orbital

What increases as you go down a group in the periodic table?

1. Atomic Radius

What decreases as you go down a group in the periodic table?

1. Electronegativity 2. Ionization Energy

What increases left to right across a period in the periodic table?

1. Ionization energy 2. Effective Nuclear charge 3. Electron Affinity 4. Electronegativity (and up and to the right)

How to determine which compound has bigger lattice energy?

1. look at charges based on if it is a metal (cations, lose electrons) or nonmetal (anions-gain electrons) 2. apply charges to coulomb's law to determine which would have a greater force

how to compare lattice energy when a cation or anion are not the same

1. multiply charges together of cation and ions 2. bigger charge = bigger lattice energy. charge and lattice energy are proportional

How many joules are in a kj?

1000 J = 1 Kj

Anion

A negatively charged ion

Cation

A positively charged ion

Compound

A substance made up of atoms of two or more different elements joined by chemical bonds

lattice

A three-dimensional array of oppositely charged ions

Why does atomic radius decrease across a period?

Across a period, effective nuclear charge increases as electron shielding remains constant. A higher effective nuclear charge causes greater attractions to the electrons, pulling the electron cloud closer to the nucleus which results in a smaller atomic radius.

Resonance structures

Two or more equally valid Lewis structures for a single molecule that differ only in the position of electrons

What are typically covalent bonds?

Typically a nonmetal bonded to a nonmetal is a covalent bond.

oxidation state (oxidation number)

a means of keeping track of how many electrons an atom has google definition: a concept that provides a way to keep track of electrons in oxidation-reduction reactions according to certain rules

the magnitude of lattice energy is....

a measure of an ionic compound's stability.

Electronegativity

a measure of how much an atom within a bond wants a molecule google definition: the ability of an atom to attract electrons when the atom is in a compound

lewis structure

a representation of covalent bonding

bond breaking =

absorbs energy

groups 1 and 2 of periodic table are...

alkaline metals (group 1) alkaline earth metals (group 2)

gravitational forces

always attractive. between two masses over a given distance

Nonmetals are usually

anions (negatively charged ions); they like to gain electrons

ion

atom that is not neutral, it has lost or gained an electron official definition: a chemical species with a net charge

In a single bond....

atoms are held together by one electron pair

In a double bond...

atoms share two pairs of electrons

According to the octet rule.....

atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration *Only valence electrons contribute to bonding

Opposite charges

attract

electromagnetic forces

attractive or repulsive forces that act between either electrically charged or magnetic object class definition: both attractive and repulsive. The force between charges (Coloumb's Law) and the magnetic force

inverse relationship with size and lattice energy...

bigger size = weaker lattice energy

positive lattice energy =

breaking bonds (absorbs energy)

Formal charge

can be used to determine the most plausible Lewis Structure when more than one possibility exists for a compound Formal charge = valence electrons - associated electrons To determine associated electrons

Metals are usually

cations (positively charged ions); they like to lose electrons

Trends for Ionic Charge

cations= positive, metals (below the stairs, and first two groups) anions= negative, nonmetals (above the stairs)

For ionic compounds, the empirical formula is also the....

chemical formula

For atomic radius, if the number of electrons is the same....

compare the number of protons.... The least number of protons = smallest atomic radii More protons = bigger atomic radii

the most effective electrons at shielding are....

core electrons

the shared pair of electrons constitutes a

covalent bond

electron affinity

deals with lone species; ability of a lone atom or ion to attract an electron.

bond length

defined as the distance between the nuclei of two covalently bonded atoms; best distance between two nuclei. Optimized or best separation between two nuclei google definition: the distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms

the chemical formula of an ionic compound

denotes the constituent elements and the ratio in which they combine

covalent bonding is....

electron sharing

In polar covalent bonds....

electrons are shared but not shared equally

pure (nonpolar) covalent bond

electrons shared equally

polar covalent bond pt. 2

electrons shared unequally

ionic bond pt. 2

electrons transferred

potential energy

energy an object possesses by virtue of its position P.E. = mass * g * h

P.E. and r are....

exponentially related. Only difference between f= Q1Q2/R^2 and P.E.=q1q2/r is the P.E. formula does not have a squared r

negative lattice energy =

forming bonds (releases energy)

reduction

gain of electrons (addition of hydrogen)

Why does atomic radius increase down a group?

in short: increased number of energy levels or principal quantum number (n) Long explanation: the number of energy levels (n) increases, so there is a greater distance between the nucleus and the outermost orbital. This results in a larger atomic radius.

as Zeff increases, ionization energy and electron affinity (the measure of an element's ability to attract an electron)....

increases for electron affinity, it's because.... it's easier to add an electron to the same shell as the positive charge of the nucleus increases.

How to compare atomic radiuses (atomic radius trened)

increases down a group, decreases across a period (left to right)

Z_eff x from left to right across the period table

increases steadily from left to right because the core electrons remain the same but Z increases for adjacent atoms

binary ionic compound

ionic compound consisting of two elements

Electronegativity pt. 2

is the ability of an atom in a compound to draw electrons to itself (not the same as electron affinity)

delta δ

is used to denote partial charges on the atoms

removing a core electron requires significantly more energy than removing a valence electron, which results in a....x.... because....

jump in ionization energy.... because: -Core electrons are closer to nucleus -Core electrons experience greater Zeff because of fewer filled shells shielding them from the nucleus.

Effective Nuclear Charge (Zeff) increases...

left to right across a period Z_eff generally changes very little down a column

When comparing lattice energies, if the charges are the same.....

look at the sizes of the ions. smallest atomic radius= highest latic energy

oxidation

loss of electrons (addition of oxygen)

empirical formula

lowest whole number ratio of the elements in a compound google definition: a chemical formula showing the ratio of elements in a compound rather than the total number of atoms

below "stairs" on periodic table

metals

Transitional metals

middle of the periodic table (elements in groups 3-12), any elements in the d-orbital portion of the periodic table -for transitional metals, take away the highest n-level e-

Compound vs. Molecule

molecule= any atoms combined even if it's the same compound= you need at least 2 different atoms. Both a molecule and a compound For compounds you need at least 2 different; thus, molecules can be composed of the same atoms but compounds cant. So not all molecules are compounds, but all compounds are molecules.

all ionic compounds are

neutral overall

Above "stairs" on periodic table

nonmental

A triple bond...

occurs when atoms are held together by three electron pairs

When force=0....

optimized separation is best. Two nuclei are at best position

Why does ionization energy decrease down a group?

outer electrons are farther from the nucleus and easier to remove. google explanation: increase in distance from nucleus and in shielding effect. This is clear because when you descend down a group, a new principal quantum shell is occupied by valence electrons. This increase in distance from nucleus and shielding effect, outweigh the increase in nuclear charge.

shielding

partial obstruction of nuclear charge by core electrons

electrostatic force can be attractive or repulsive. repulsive is.... attractive is....

positive sign = repulsive force (think -1 x -1 for the charge of two electrons) negative sign = attractive force (think -1 x 1 for an electron and a proton) *both attractive and repulsive forces are present in atoms

the formation of ionic bonds....

releases a large amount of energy

bond made =

releases energy

Like charges

repel

When compounds form between elements with similar properties, electrons are not transferred from one element to another but instead are...x.... in order to give each atom a noble gas configuration.

shared

Shared electron pairs

shown either as dashes or as pairs of dots

The shorter multiple bonds are stronger than....

single bonds

multiple bonds are shorter than...

single bonds

atomic radius

size of an atom

In a series of isoelectronic ions, cations are *x* than anions.

smaller

reducing agent (reductant)

species that is oxidized google definition: The substance that is oxidized and thereby causes the reduction of some other substance in an oxidation-reduction reaction.

oxidizing agent (oxidant)

species that is reduced google definition: the substance that is reduced and thereby causes the oxidation of some other substance in an oxidation-reduction reaction.

effective nuclear charge (Zeff)

the actual magnitude of positive charge that is "experienced" by an electron in the atom Zeff = Z - σ Z= number of protons/atomic number σ= number of core electrons

lattice energy

the amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase

Electro static force

the attractive or repulsive force between two electrically charged objects

ionization energy

the energy required to remove an electron from an atom official definition: the least amount of energy required to completely remove an electron from 1 MOLE of that gaseous species, when it is in its ground state

Central atom in a Lewis Structure

the least electronegative that isn't hydrogen

Lattice energy depends on.....

the magnitudes of the charge and on the distance between them

orbital penetration

the proximity to which an electron can approach the nucleus

Lewis theory of bonding

when compounds form between elements with similar properties, electrons are not transferred from one element to another but instead are shared in order to give each atom a noble gas configuration

covalent bonding

when two atoms share a pair of electrons