Ionization and Ionic Bonds

Ionization Energy: Period

Ionization energy increases across a period. Zeff increases but shielding is constant. So is distance. Harder to remove an electron.

Using Electron Configuration

Beyond period 3 ignore numbers that aren't the largest "n" value. Unless discussing transition or inner transition metals only count s and p orbitals.

Henry Mosely

Defined the theory of atomic number. Determined the atomic number of each element. Ordering by atomic number fixed all the problems with Mendeleev's table.

Coulomb's Law

Describes the electrostatic force felt between charged particles Fe ∝ (q1 • q2) /r^2. Fe is electrostatic repulsive force. q are the respective charges. r is distance between charges.

d - block

Across period adding electrons to d-sublevel - non valence electrons. Nuclear charge up but so is shielding. Very little change.

Naming Anions

Anion do not bear same name as neutral element. Swap out end of word for -ide.

Sea of Electrons

As opposed to all other types of bonding, metals are most accurately modeled as a "fluid" substance. Valence electrons do not belong to any single nucleus. Metal nuclei in a sea of electrons.

Coulomb's Law: Distance

Assuming the magnitude and charge of each remains the same, the force between the particles will decrease quickly as the distance between them increases.

Radius: Groups

Atomic radius increases going down a group. With each new period - principle quantum number increase. Each subsequent energy level is farther away from the nucleus: r increase: electrostatic force is lower. Atomic radius increases as energy level does. Also because of shielding.

Transition Metals

Can form from multiple different cations. Have to look up which cations form. Don´t need to have noble gas electron configurations. An octet rule exception.

Ionization Equation

Can write a chemical equation for this process. Base metal on the left. Ion and electrons on right. Electron on left if negative ion. Conservation of Charge.

Electronegativity

Chemical property describing elements' tendency to attract electrons in a chemical bond. Every element is assigned a number based on attractive force (important in covalent bonding). As Zeff increases so does electronegativity. As distance increases Zeff decreases. Increases across period. Decreases down group. No value assigned for noble gases, stable.

Ionic Compounds

Composed of cation(s) and anion(s). Charged ions but compounds are electronically neutral. Any positive charge must be canceled out be a negative one.

Metal Implications

Conductivity: electrons carry electric charge. Electrons flow fight through the sea. Also able to be pushed ductile and malleable.

Valence Electrons

Electrons in the outermost shell (principle quantum number) that will react and interact with other atoms. Periods 1-3 giver the electron configuration when abbreviated. This allows for the rationalization that elements in the same groups have similar chemical properties.

Reactivity on the Periodic Table

Elements in the same group have similar chemical and physical properties. Electrons and electron configuration determines reactivity.

Electron Transfer

Elements ionize to ideal charge via electron transfer. Transferred from metal to nonmetal. Electrostatic attraction keep them together.

Coulomb's Law: Atoms

Fe ∝ (q nuc •q e) / r^2. q1 is the nucleus q2 is any given electron and r is the distance between them. Nucleus is tightly packed and treated as a unit; electrons are all separate and must be analyzed individually.

Number of Valence Electrons

Groups give number of valence electrons. Groups 13-18 subtract 10.

Halides

Halogens all form 1-. Ions called halides.

Properties of Ionic Crystals

Highly ordered and extremely stable. Have high melting points. Soluble in water. Salts do not conduct electricity. Solutions in water are highly conductive - freely moving ions carry charge. Most ionic compounds are solid crystals. Highly ordered 3 dimensional repeating units. Several "common" crystal structures. Contains of unit cell and repeats in three dimensions,

Alloys

Homogeneous mixture of two or more metals or solutions. Alloys combine the properties of their constituent metals. Properties are often superior to those of their components. Can tailor to needs.

Ionization Energy: Groups

Ionization energy decreases down a group because it is easier to remove an electron. With a new energy level, the shielding increases but Zeff is the same. Highest occupied energy level is farther so r increases.

Coulomb's Law: Charge

If both charges are + or - the repulsive force is positive. If they are opposites, negative attraction force. If magnitude of charge increases, magnitude of force increases.

Multiple Ionizations

If multiple electrons are removed each has required energy associated with it. Further ionization require more energy (protons > electrons). Once it hits noble gas configuration hard to remove an electron because it is jumping down principle quantum number. Large jump in ionization energy.

Atomic Radius

In atomic orbitals, edges were fuzzy. To measure atomic radius we find the distance between two nuclei in pico meters (10^-12). One might expect atomic radius to increase with each electrons. Radius increases down a group but decreases across a period.

Ionic Bonding

Ions are held together by electrostatic force. +/- attraction. Ionic compounds are neutral.

Organization of the Periodic Table

It became intuitive to categorize the elements by reactivity and chemical properties. In the beginning of the 19th C. many variations of a table, but none widely accepted.

Metals

Left of the metalloids. 80% of elements are metals. Good conductors of heat and electricity. Lustrous when freshly cut. Reflect light. Ductile (wires) and malleable (sheets). Solid at room temp - except Hg.

Lewis Dot Structure

Lewis dot structures help us quickly and clearly represent valence electrons. Draw the atomic symbols. Count number of valence electron. Add dots for each electron starting with one electron one each side of the symbol. Then pair up counterclockwise.

Formation of Cation

Main group metals form cations by losing valence electrons, leaving a complete octet at the next lowest energy level. The noble gas that is in brackets.

Metalloids

Metal-like and nonmetal-like properties. Properties can be manipulated under different conditions. Sulfur is a bad conductor but with a small amount of Boron becomes a good conductor.

Formation of Anions

Nonmetals form anions by gaining electrons to form a new octet - configuration of the next noble gas. Atoms will predictably gain the number of electrons needed to make eight.

Simple Cubic

One lattice point (atom/ion) on each corner of a cube. Total one atom per unit cell (1/8 x 8). Relatively rare (ex polonium).

Body Centered Cubic

One lattice point on each corner and one in the center of the unit cell. Total two atoms per unit cell. ex CsCl.

Face Centered Cubic

One lattice point on each corner of the cube and lattice points on the faces of a cube that give 1/2 atom contribution. Total four atoms per unit cell. ex NaCl.

Ion properties

Properties of the ion can be vastly different from that of atoms. Na+ has more in common with a noble gas - nonreactive. Cation bears the same name as atom ____ ion.

Reactivity: Nonmetals

React by combining or gaining electrons. Reactive nonmetals have high ionization energy. Reactivity decreases down group and increases across period.

Reactivity: Metals

React by losing electrons. More reactive metals have lower ionization energy. Reactivity increases down group and decreases across period.

Example of Alloys

Silver vs Sterling Silver: Sterling silver is cheaper and stronger. Used for utensils. 92.5% Ag 7.5% Cu. Bronze: more durable than copper and easily cast. 87.5% copper 12.5% tin. Steel: contain iron and chromium. Others for specific uses: stainless steel, surgical steel. Cast Iron - conductive - non-malleable. Golds

Crystal Structures of Metals

Stable, closely packed structure. Similar to ionic crystal structures. All the same atom. Talking about the nuclei.

Types of Alloys

Substitutional Alloy: added metal takes place of a base metal. Must be of reasonably same size and properties. Interstitial Alloy: added metal occupies the spaces between base metal and must be smaller than base metal.

Octet

The most stable form of an atom that will have eight valence electrons (full s & p sublevels) - noble gases represent a full octet. "Octet rule".

Ionization Energy

The energy required to remoce one electron from an atom. Always measured in the gas state. First electron removed - cation. The larger the force exerted on the electron from the nucleus the harder it will be to remove that electron - high ionization energy.

Periodic Table

The organization of all known elements. Only a few elements were known in the 18th century so there was no need to organize them. As are were discovered there was a upshot be able to group and organize them. Lots of information on the table with a minimum of the atomic symbols, number, and mass.

Periodic Predictions

The periodic table can be used to predict many properties including size, ionization, ionization energy, ion size, and electronegativity.

Force of Attraction

The radius of an atom is determined by the force of attraction felt between the charged nucleus and the outer most electrons. If force of attraction is greater, will pull electrons closer to the nucleus.

Bravais Lattices

Three main cubic crystal structures. Simple cubic, body centered cubic, face centered cubic. This is for crystallized elemental solids and ionic solids.

Nonmetals

To the right of the metalloids. More variation in properties. Mostly gases with some solids and one liquid. They are brittle and poor conductors. Carbon is an exception.

Radius: Periods

To understand this trend we must develop understanding of effective nuclear charge (Zeff). Nuclear charge (Z) is sum of charge of the protons - atomic number. Effective nuclear charge - total positive charge felt be the electron. Moving across period Z increases but so does Zeff. Fe ∝ (Zeff • q e) / r^2. Zeff increases nucleus exerts a stronger force. Radius gets smaller.

Force of Repulsion

What keeps the electrons from being pulled all the way to the nucleus. Fe ∝ (q e • q e) / r^2. Between two electrons both negatively charged making repulsive force positive. Balance the inward pull of nucleus and outward push of electrons.

Ions

When atoms gain an electron, the electrons outnumber the protons - anion - negatively charged. When atoms lose an electron, protons outnumber the electrons - cation - positive.

Transition / Inner Transition Metals

When comparing elements in d &f blocks only compare the electrons in that sub level. More complicated than just valence electrons.

Ionic Size

When electrons are added, electron electron repulsion increases and electron cloud expands to compensate - radius increases. When electrons released electron electron repulsion decreases and electron cloud contracts - radius decreases. The greater the magnitude of the charge the more exaggerated. Effects of charge can out weigh group trend.

Metal Ionization

When metals ionize they tend to lose electrons from their highest occupied energy level. Form a cation.

Nonmetals Ionization

When nonmetals ionize they tend to gain electrons into their lowest unoccupied energy level, Form an anion.

Z vs Zeff

Why are these values different? Outer electron is shielded from nucleus by electrons in between - cancels some charge. Can be calculated by Zeff = Z - S. Z is atomic number, S is the average number of electrons between the nucleus and the electron in question. Remains constant across a period.

Mendeleev's Table

Wrote the properties on notecards and then arranged elements until they made sense. Ended up ordering them by atomic mass. Strengths: It was able to make sense of shared properties as part of repeating trends. Was able to predict elements that had not been discovered by aligning elements with know properties. Weaknesses: Because of ordering by atomic mass brought up a few issues. For example: iodine has a larger atomic number but smaller atomic mass than Te. If it was ordered by atomic mass, Iodine would not be with the halogens. Mendeleev said that the masses were wrong but they were not.