DBA Notes For Module 3

Lesson 03.01: History of the Periodic Table
Explain how scientific observations led to the development of, and changes to, the periodic table.
Russian chemist Dmitri Mendeleev set out to organize the 63 known elements according to their properties.
Mendeleev wrote the name, mass, and chemical properties of each element on a separate card and arranged the cards by various properties, looking for trends and patterns.
When he arranged the elements by increasing atomic mass, Mendeleev noticed that similar chemical properties repeated at regular intervals. This type of repeated pattern is called “periodic,” which is where the periodic table gets its name.
The first question, about the out-of-order elements, was not answered until 40 years later. The English scientist Henry Gwyn Jeffreys Moseley, who was working with Ernest Rutherford, performed a series of experiments with 38 metals.
As Moseley analyzed these elements, he found that the positive charge of each element’s nucleus increased by one from element to element as they were arranged in Mendeleev’s periodic table.
Moseley’s findings led to: the modern definition of atomic number: the number of protons in an atom’s nucleus recognition that atomic number, not atomic mass, is the basis for the organization of the periodic table
Although Moseley’s findings changed the basis of periodicity from atomic mass to atomic number, they actually justified Mendeleev’s organization of elements on his periodic table.
Moseley’s atomic number supported and justified the observed periodic properties of elements and the organization of elements on Mendeleev’s periodic table.
Describe the organization of the modern periodic table.
In the modern periodic table, elements are arranged from left to right in rows, called periods, by increasing atomic number, and from top to bottom in columns, called groups, based on similar chemical properties.
Lesson 03.02: Group Names and Properties
Compare and contrast the properties of metals, metalloids, and nonmetals.
Metals
are good conductors of heat and electricity reflect light and heat have a luster malleable, meaning they can be hammered or rolled into thin sheets ductile, meaning they can be pulled into thin wires without breaking
Metals can be used for construction, in electrical devices, and in many other uses.
Non-Metals
poor conductors of heat & electricity oxygen, hydrogen, chlorine, fluorine, and nitrogen are gases at room temperature are brittle
Metalloids
semiconductor that has some characteristics of metals and other characteristics more like nonmetals. silicon and germanium
Identify groups and sections of the periodic table by group name and common properties.
Alkali Metals
The elements in Group 1 of the periodic table are known as the alkali metals. These metals are extremely reactive, reacting with water and with most nonmetals.
All of the alkali metals have a silvery appearance and they are soft enough to be cut with a knife.
Alkali metals are so reactive that they do not occur in nature as pure elements; they are always found in compounds. Many of these compounds are essential to our lives, including being found in the fluids that fill and surround all of our cells.
The metals of this group all have one valence electron, which they give up when they react to form ionic compounds.
Alkaline Earth Metals the elements of Group 2 are called the alkaline-earth metals. These metals share similar properties because they each have two electrons in the s sublevel of their outermost energy level. Metals in Group 2 are harder, denser, and stronger than the metals in Group 1.

Although they are less reactive than the Group 1 alkali metals, the alkaline-earth metals are still too reactive to be found as pure elements in nature. Compounds containing alkaline-earth metals are found in the earth’s crust, in sea salts, and in gems and minerals like emerald and aquamarine.
Halogens
The elements in Group 17 are known as the halogens. The halogens are the most reactive of the nonmetals; they react with most metals to form ionic compounds known as salts. As pure elements, the halogens are diatomic molecules.

This means that the halogens bond in pairs, forming compounds such as F2, Cl2, and Br2. As you will learn, the high reactivity of the halogens is based on the presence of seven valence electrons, one electron short of having a full valence. Some of the halogens (F2 and Cl2) are gases at room temperature, while others are liquids (Br2) or solids (I2). The halogens are used in water purification, photography, insecticides, bleaches, plastics, and many other practical applications.
Noble Gases
Group 18, the group at the far right of the periodic table, contains nonmetals called the noble gases. The first noble gas to be discovered on Earth was argon, which was discovered in 1894 by the English physicists John Williams Strutt and William Ramsay.

Argon makes up one percent of Earth’s atmosphere, but it escaped notice for so long because it does not react with any other elements. Another noble gas, helium, was discovered earlier in the sun’s light spectrum, but Ramsay was the first to discover that helium is also found on Earth. The noble gases are not found in compounds in nature because their full valence energy levels lead to very low reactivity. All noble gases have a full valence; helium has two electrons in the first energy level, while all other noble gases have eight valence electrons.

Transition Metals
The metals in the d-block, groups 3–12, of the periodic table are called transition metals or transition elements. They are good conductors of electricity and are typically less reactive than the metals in groups 1 and 2. Many of these metals can be found as pure elements in nature, and are used in construction, electrical wiring, and even jewelry. You will find that you may be familiar with many of the transition metals, such as silver, gold, copper, iron, mercury, and platinum.

Transition Metals
The lanthanides, the 14 elements found between lanthanum (La) and hafnium (Hf) in Period 6, are all shiny and reactive metals. Chemists use modern separation techniques to isolate these elements to be used in practical applications, like inside television tubes.

The 14 elements found between actinium (Ac) and rutherfordium (Rf) in Period 7 are called the actinides. The first four actinides can be found naturally on Earth, while the rest have only been synthesized in a laboratory. The actinides are unstable and radioactive, so these elements are not used very often in everyday chemistry.
Lesson 03.03: Periodic Trends
Describe and explain the trends for effective nuclear charge, atomic radius, ionic radius, and ionization energy across a period and down a group.

The atomic radius of an element is half the distance between the centers of two atoms of that element that are bonded together.

There is a gradual decrease in atomic radii from left to right across a period. The decrease in radius is due to the additional number of protons in each nucleus.

There is a general increase in atomic radii going down each group of elements on the periodic table. Going down the group, there is an increase in the number of occupied energy levels in the electron cloud, as well as an increase in protons.

The effective nuclear charge felt by an atom’s valence electrons increases by one for each element from left to right in a period. This is because each element has one additional proton in its nucleus, adding to the overall nuclear charge.

The effective nuclear charge felt by an atom’s valence electrons stays constant for each element down a group. This is because the number of additional protons in each element’s nucleus is equal to the number of additional shielding electrons in the inner energy levels of the atom.

Effective Nuclear Charge is the charge (from the nucleus) felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus
Subtracting the number of shielding (core) electrons from the total nuclear charge (number of protons) gives the effective nuclear charge felt by an atom’s valence electrons.

Ionization Energy is the energy required to remove one electron from an element, resulting in a positive ion.
Ionization energy has a general increase for elements across a period from left to right. The stronger the effective nuclear charge felt by an atom’s valence electrons, the more energy is required to remove one of those electrons from the atom.
The ionization energy of elements decreases going down a group because the atomic radius of the atoms increases.

An ionic radius is one-half the diameter of an ion.

A positive ion is called a cation, and a negative ion is called an anion.

There is a decrease in the ionic radii of the cations from left to right across the metals, and also in the ionic radii of the anions from left to right across the nonmetals.

Ionic radii increase down a group, following the same trend as atomic radii, because there is still an increase in the number of occupied energy levels in the electron cloud.
Lesson 03.04: Valence Electrons and Bonding
Define and compare ionic and covalent bonding.
Ionic Bond: A chemical bond that results from electrostatic attraction between positive and negative ions.
Covalent Bond: A chemical bond where electrons are shared between two atoms, neither atom completely gains or loses electrons.
Relate your knowledge of the periodic trends to the chemical bonding exhibited by various elements.
Atoms are held together in compounds by electrostatic attraction between positive nuclei and negative electrons. This attraction holds atoms together in a chemical bond, a link between two atoms resulting from the mutual attraction of their nuclei for valence electrons. All chemical bonds involve valence electrons, but the bonds are classified by the way in which electrons are distributed within the bonds.
Lesson 03.05: Ionic Bonding and Writing Formulas
Determine an element’s ionic charge based on its location on the periodic table.
Elements in the same group of the periodic table have the same number of electrons in their valence energy levels, so they tend to have the same charge when they form naturally occurring ions. Just as the periodic table was your guide for electron configurations and valence electrons, it also helps you determine the ionic charges of the main group elements.

As you know, opposite charges attract. The attraction between positive and negative ions is called an ionic bond. Even though ions are charged, they combine in ratios that cancel those charges and make the ionic compound neutral. When you write the formula of an ionic compound, subscripts are used to represent the ratio of positive and negative ions in the ionic compound

Write the correct ionic formula when given two elements that bond ionically.
The formula of an ionic compound does not represent an individual unit that can be isolated or exist individually. Instead, it represents a ratio of how the positive ions and negative ions combine in an ionic crystal. The formula of an ionic compound is called a formula unit, which represents the simplest whole number ratio of positive to negative ions in the ionic compound.

Lesson 03.06: Covalent Bonding and Lewis Structures
Determine how many covalent bonds an atom needs in order to fill its valence shell, using the periodic table.
The number of covalent bonds an atom can form is usually equal to the number of electrons it needs to fill its valence energy level. The halogens in Group 17 have seven valence electrons and need one more electron to fill their valence energy level, so they usually form one covalent bond. Oxygen and sulfur have six valence electrons and need two more electrons to fill their valence energy level, so they usually form two covalent bonds when bonding with other nonmetals. Hydrogen has one valence electron and need one more electron to fill their valence energy level, so they usually form one covalent bond.
Draw correct Lewis structures to model covalently bonded molecules when given the name or formula of the molecule.
1. Place the least electronegative element in the center of the molecule. The central atom is usually the first element written in the molecular formula, except H, which cannot be a central atom because it only bonds once and then its valence is “full.”
2. Calculate the total number of valence electrons. Add up the total number of valence electrons that should be in the picture according to the periodic table. Make a list of the number of valence electrons (determined by the location on the periodic table) for each atom in the molecule, and then add them together. We will call this the “reality number” because this is the number of valence electrons that must be present in the molecule.
3. Write the skeleton structure. Attach all other atoms to the center with single bonds. You may change this later, but we know that each atom in the molecule must be attached with at least one set of shared electrons (a single covalent bond).
4. Complete the valence electrons. Fill every atom’s valence by adding electron pairs until they are all full. Remember that most atoms need eight valence electrons to be full, except H, which is full with two electrons. Also, remember that a single bond (each line drawn in the model) represents two electrons being shared by both atoms involved in the bond. This means that one line counts as two electrons in the valence of each atom touching that line.
5. Tally the totals. Count up the number of electrons represented in your drawing. Remember that each line represents two shared electrons and each pair of dots represents two unshared electrons when you are counting the electrons in the drawing. We will call this counted number the “picture number” in our examples.
6. Perform a comparison. Compare the picture number to the reality number to see if you need to change your Lewis structure drawing.
a. If picture # = Reality #, then the Lewis structure drawing is a good representation of the molecule and you do not need to change anything.
b. If picture # > Reality #, you must fix it by adding multiple bonds where appropriate in order to reduce the Picture # to equal Reality #.
When adding an extra bond, erase a lone pair from each atom touching that bond. That means that for each additional bond drawn, the total of the picture number is reduced by two.
Add the extra bond where allowed (make sure that the atoms have room in their valence energy levels to form multiple bonds, and try to keep the molecule symmetrical whenever possible).
7. Distribute. Distribute electrons to atoms surrounding the central atom to satisfy octet rule.** If less than eight electrons on central atom, it means a multiple bond: two electrons less means double bond; four electrons less means triple bond.
a. Atoms that form multiple bonds are C, N, O, S. Oxygen atoms do not bond to each other (except in O2 & O3, H2O2, peroxides, superoxides).
b. All atoms must have eight electrons (octet rule), except hydrogen whose octet is two.
Describe your observations and conclusions from the virtual lab.
A and C are Ionic Bonds because their melting points were over 300C, and their liquid states had conductive properties. Substances B and D are Covalent Bonds because they are under 300C and are not conductive in any form.

Ionic Compounds have higher melting and boiling points than covalent compounds.
Do covalent compounds conduct electricity as: (3 points)
Solids? No
Liquids? No
Aqueous solutions (when the covalent compounds are dissolved in water)? No

Do ionic compounds conduct electricity as: (3 points)
Solids? No they are insulators
Liquids? Yes
Aqueous solutions (when the ionic compounds are dissolved in water)? Yes
Lesson 03.07: Intermolecular Forces
Use VSEPR theory to predict the shape of a molecule based on its Lewis structure.
The method we will be using to predict the shape, or geometry, of a molecule is called the valence shell electron pair repulsion (VSEPR) theory. The VSEPR theory is about geometry of compounds and electron location. Electron pairs in a molecule repel one another; therefore electrons will distribute themselves in positions around the central atom that are as far away from each other as possible. The VSEPR shape of a molecule can be determined after drawing a Lewis dot structure.
In chemistry, polarity refers to a separation of electric charge leading a molecule to have a dipole moment, a net electrical charge. The molecular polarity is dependent on the difference in electronegativity between atoms in a compound and the asymmetry of the compound's structure. Polarity is a physical property of compounds which relates other physical properties such as melting and boiling points, solubility, and intermolecular interactions between molecules.
Permanent dipoles: These occur when two atoms in a molecule have substantially different electronegativity: A molecule with a permanent dipole moment is called a polar molecule.
Instantaneous dipoles: These occur due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a temporary dipole. A molecule is polarized when it carries an instantaneous or an induced dipole.
Induced dipoles: These can occur when one molecule with a permanent dipole repels another molecule's electrons, "inducing" a dipole moment in that molecule temporarily.
Compare and contrast intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole).
London dispersion forces occur between all molecules and particles but are the only force of attraction between nonpolar molecules or noble gas atoms. These forces are the weakest of the intermolecular forces.
Dipole-dipole forces are electrostatic interactions of permanent dipoles in polar molecules. The attractive forces that occur between the positive end of one polar molecule and the negative end of another polar molecule tend to align the molecules to increase the attraction.
Hydrogen bonding is a particularly strong dipole-dipole interaction in which hydrogen is covalently bonded to a highly electronegative element, and attracted to the very electronegative element in another molecule
Ion-dipole forces are attractive forces that result from the electrostatic attraction between an ionic compound and a polar molecule. This interaction is most commonly found in solutions, especially in solutions of ionic compounds in polar solvents, such as water.
Identify the intermolecular forces experienced by different compounds.
Intermolecular forces are attractive forces between molecules. These forces exist between molecules when they are sufficiently close to each other. They are responsible for the non-ideal behavior of gases and for properties of matter such as boiling point and melting point. In contrast to intramolecular forces, intermolecular forces are much weaker.

Lesson 03.08: Naming Compounds
Correctly name covalent compounds, ionic compounds, and acids when given their formulas.
1. Look at the first element in the formula. Ca is the symbol for Calcium. The names of the elements are found on the periodic table. The first word in the binary ionic compound name is Calcium.
2. Take the name of the nonmetal. Cl would be the nonmetal because it is second in the chemical formula. It is the symbol for the element Chlorine. Shorten Chlorine to its root word Chlor-.
3. Add –ide to the root word. In this case, Chlor- would add with –ide to form Chloride.
Lesson 03.09: Molar Mass of Compounds
Calculate the molar mass of compounds from the formula.
Because a formula represents the amount of each type of element within a compound, the molar mass of a compound can be determined using the molar masses of each element multiplied by that element’s subscript within the formula.
(6.022 × 1023 atoms)
Determine empirical formulas from percent by mass or mass data.
A formula for a molecule that is expressed as the smallest possible whole-number ratio of the elements in the compound is called an empirical formula.
1. Convert the amount of each element from grams to moles by using their molar mass values from the periodic table.
2. Divide each of the mole values by the lowest number in order to get a whole-number mole ratio. The mole values are a mole ratio, but an empirical formula’s subscripts must be the lowest whole-number ratio between the elements.
3. After dividing, most or all of the mole values should be whole numbers. If not, you must multiply all of the mole values by a number that would make them all whole numbers.
If any of the mole values ends in .5 you must multiply all of the mole values by 2 in order to get all whole numbers.
If any of the mole values end in .33, you must multiply all of the mole values by 3 in order to get a ratio of all whole numbers.

Determine the molecular formula from the empirical formula and molar mass of a substance.
1. Calculate the mass of the empirical formula in the same way you would calculate the molar mass of a compound.
2. Divide the molecular mass (the actual molar mass of the compound) by the mass of the empirical formula in order to find the ratio between the molecular formula and the empirical formula. The answer to the division should be a whole number.
3. Multiply all subscripts in the empirical formula by this ratio to give you the molecular formula.