Introduction to General Chemistry

All mass consists of tiny particles called atoms. Each atom is composed of a nucleus surrounded by one or more _____. The _____ of a nucleus is on the order of 10-4 angstroms (A). One angstrom is 10-10 m. The nucleus consists of _____ and ______, collectively called _____.

Protons and neutrons are approximately equal in size and mass. Protons have a positive charge, while neutrons are electrically neutral. Protons and neutrons are held together to form the nucleus by the _____ _____ ____. The stability of the nucleus can be measured by its _____ _____ (define)

Surrounding the nucleus at a distance of about 1 to 3 A are electrons. The mass of an electron is more than _____ times smaller than the mass of a nucleon. Since the nucleons are ____ compared to the distance between the nucleus and the outermost electrons, the atom itself is composed mostly of ____ ____. If an atom were the size of a modern football stadium, it would have a nucleus the size of a _____. Matter is composed of atoms, therefore matter is mostly _____ ____.

State which of the three atom constituents is designated and state its charge

Electrons and protons have _____ charges of _____ magnitude. Although the charge on an electron is written as 1- and the charge on a proton is written as 1 +, remember that this charge is in electron units 'e' called the _____ ______. A charge of 1 e is equal to _______ coulombs, the SI unit for _____. When is an atom is electrically neutral?

Elements are the building blocks of ______ and cannot be decomposed into simpler substances by _____ means. An element may have any number of ______ or ______, but only one number of _____.

All elements can be depicted as shown in Figure 1.2. Z is the _____ number, indicating the number of protons. The atomic number provides the identity of the element because each element ,has a unique number of _____: if we know the atomic number, we know the _____ of the element.
Label:

The mass number, A, is the number of ____ plus _____, and varies depending on the number of _____. Protons and neutrons each have a mass of approximately ___ amu and the mass of an atom is concentrated in the _____. This means that the mass number of an element is approximately equal to its ____ ____ or _____ _____.

The atomic weight given on the periodic table can be assigned either of the commonly used units: _____ _____ _____ (abbreviated as amu or the less commonly used SI abbreviation, u) or ____/____.

Why is it that not every atom of a given element will have a mass equal to the mass number? What is the atomic weight is equal to?

The nucleus of a specific isotope (define) is called a ____. Isotopes have similar _______ properties. Hydrogen has three important isotopes:____, ______, and _____ 99.98% of naturally occurring hydrogen is _____.

The isotopes for carbon include ____, ____, and ____. Each of carbon's isotopes contains ___ protons with _____ ____ of neutrons. The ____ protons are what define carbon; if the number of _____ changed, it would no longer be carbon. 12C (carbon-12) contains ___ neutrons, 13C (carbon-13) contains ___ neutrons , and 14C (carbon-14) contains ___ neutrons.

Ions (define) are not electrically _____. Positive ions have fewer electrons than protons and are called _____; negative ions have more electrons than protons and are called _____. Define salt

When a neutral atom loses an electron to become a cation, it gets ______. The atom still has the same number of ____, so there are now more _____ than ______. The result is that the positive charge of the nucleus exerts a greater attractive force on each _____ _____, pulling them closer to the ______.

The loss of an electron also ______ the repulsive forces between the electrons, further contributing to the _____ in size. On the other hand, when a neutral atom gains an electron to become an _____ , it gets ______. The atom now has more electrons than protons, so the positive charge of the nucleus pulls _____ strongly on each individual valence electron. The addition of an electron also ______ the repulsive forces between the electrons, pushing them farther away from each other. The net effect is that the anion is _____ than the neutral atom.

The periodic table lists the elements from left to right in the order of their _____ numbers. Each horizontal row is called a _____. The vertical columns are called _____ or _____. As will be described in more detail below, elements in the same family share some similar ______ and ______ properties.

The periodic table in Figure 1.4 divides the elements into three sections:

*There are two methods commonly used to number the groups. The newer method is to number them 1 through 18 from left to right. An older method that is still occasionally used is to separate the groups into sections A and B, and then number them with R oman numerals as show n in Figure 1.4. The periodic table on the MCAT® does not have group numbers.

Metals (define) can be described as atoms in a sea of ______, which emphasizes their loose hold on their electrons and the fluid-like nature of their _____ _____. What gives them (metals) their metallic character?

4 metallic characteristics of metals

Metal atoms easily ____ ____ each other, allowing metals to be hammered in to thin sheets or drawn into wires. _____ move easily from one metal atom to the next, transferring _____ or _____ in the form of _____ or _____. All metals except ______ exist as solids at room temperature. Metals typically lose ______ to become cations, and form _____ bonds.

Nonmetals have diverse _____ and ______ behaviors. Molecular substances are made from nonmetals, since nonmetals form _____ bonds with one another. Generally speaking, nonmetals have _____ melting points than metals. They tend to form _____, which commonly react with ____ _____ to form ionic compounds. Metalloids have some ______ and some ___-_____ characteristics.

The section A groups (groups 1, 2, 13, 14, 15, 16, 17, 18) are known as the _______ elements or main-group elements and the section B groups (groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12) are called the _____ _____.

It is important to be able to predict what kind of ion an element will form, i.e. whether it will gain or lose ______. The representative elements make ions by forming the closest _____ _____ _____ configuration; this is why metals tend to form cations and nonmetals tend to form _____.

When the transition metals form ions, they lose electrons first from their _____ ______ and then from their ______, as will be discussed later in this lecture.

The periodic table below shows some of the common ions formed by the _____ _____. Notice in the figure below that Group ___ elements make 1+ ions. There are five 3+ ions: ___, ___, ___, ___, and ___. The other transition metal ions are ___.

Easy here! You don't have to memorize the charge on every cation made by transition metals. This is background knowledge. But you should be able to predict the charge based upon two things:

Elements in the same group on the periodic table have _____ chemical properties (explain)

The MCAT® requires the ability to identify and make predictions about specific groups: Group 1 (_____ ____), Group 2 (____ _____ ____), Group 16 (the ____ group), Group 17 (_____), and Group 18 (_____ ____). *This section will consider the properties that are useful for making predictions about elements in these groups.

State 6 characteristics of the Group 1 or alkali metals as pure substances

6 characteristics of hydrogen

5 characteristics of the alkaline earth metals (group 2)

3 characteristics of group 14 elements

7 Characteristics of Group 15 elelements

9 characteristics of the Group 16 elements

8 characteristics of the group 17 elements

4 characteristics of the group 18 elements

The elements that tend to exist as diatomic molecules are _____, _____, _____, and _____. When these elements are discussed, it is safe to assume that they are in their ______ form unless otherwise stated. In other words, the statement "nitrogen is nonreactive" refers to ___ rather than ___.

This section will describe how the periodic table can be used to predict changes in the chemical properties of atoms across a period or down a group. These predictions can be summarized as four periodic trends: _____ _____, _____ _____, ________ and _____ _____. As a general rule, the atomic radius increases going "____" and "to the ____" on the periodic table, while the other properties increase "___" and "to the _____."

The atomic radius is the distance from the center of the _____ to the outermost _____. Atomic radius corresponds to the size of the atom. Moving across a period, the radius ______ (explain)

However, moving down a group, new shells of ______ are added. These outer electrons are "______" from the attraction of the _____ in the inner nucleus. As a result, atomic radius ______ going down a group.

To gain a deeper understanding of changes in atomic radius, and to understand the other periodic trends, consider Coulomb's law, F =kq₁q₂/r₂. Coulomb's law describes the _______ force (define) between an ______ and the _____.

With respect to coulomb's law, the distance between the electron and the nucleus is __. For q1 we might plug in the ______ charge of the nucleus, (__), and for q2 , the charge on an ______, (__). This would work well for hydrogen, where the lone electron feels ___% of the positive charge on the nucleus.

However, in helium the _____ electron shields some of the nuclear charge from the _____ electron, so that it doesn't feel the entire nuclear charge Z. Shielding occurs due to the repulsive forces between ______.

The amount of charge felt by the most recently added electron is called the ______ _______ ______ (___)· In complete shielding, each ______ added to an atom would be completely shielded from the attractive force of all the _____ except for the _____ _____ added, and the Zeff would be ___ eV for each electron.

Without shielding, each electron added would feel the full ______ force of all the ______ in the nucleus, and the Zeff would simply be equal to ___ for each electron.

Figure 1.6 shows Zeff values (given in electron-volts) for the highest energy electron in each element through sodium. Zeff generally _____ going left to right across the periodic table. While more protons are added across a period, ______ Z, the new electrons added are in Electron Salty roughly the same energy level, and therefore do not experience significantly more ______ than the previous electron.

Zeff also _____ going from top to bottom down the periodic table. Though the energy level of the outermost electrons _____ down a group, the attractive pull of the growing positively charged nucleus ______ the additional shielding effects of higher electron shells.

In Figure 1.6, notice that Zeff drops going from neon to sodium (explain). This causes a strong ______ in shielding and ______ in Zeff, but the outermost electron in sodium still experiences a _____ Zeff than the outermost electron of the element immediately above it on the periodic table, lithium (explain). A similar drop also occurred between He and Li, though it was not quite as large because there were fewer _____ and _____ involved.

The ____, and not __, should be plugged in for q1 in Coulomb's law to find the force on the outermost electron. The force on an electron is a function of both q1 (___) and r (the distance from the ____). Notice that Zeff can be used to explain the atomic radius trend discussed earlier. (Demonstrate: 3-Story)

Zeff can also be used to understand trends in ion size. As discussed previously, cations are _____ than the neutral element while anions are _____. When an atom loses an electron, Zeff _____ because there are now more _____ relative to _____. ______Zeff means electrons are pulled closer to the nucleus. When an atom gains an electron, Zeff ______ because there are now more ______ relative to _____. Shielding increases due to the increased ______ _____ between electrons, and ______ are pushed further away from the nucleus.

Isoelectronic ions are ions with the same number of ______, but different ______ identities. O^2-, F, neutral Ne, Na+, and Mg2+ all have the same number of electrons. However, since they have different numbers of protons, the electrons feel different ______ ______ _____. As a result, isoelectronic ions are not the same ____. The largest of these is ___. Its electrons feel the ______ Zeff (why?). The smallest of these is Mg2+. Its electrons feel the ______ Zeff (explain)

Ionization energy (define) generally increases along the periodic table from ____ to _____ and from _____ to ____. When an electron is more strongly attracted to the nucleus, _____ energy is required to detach it.

Define first and second ionization energy

The second ionization energy is always _____ than the first (explain)

For periodic trends, all the "E's" are up and to the right: (list them)
What happens to the other trend

The ionization energy trend can be remembered by considering its relation to
zeff and r. Moving across a period to the right, ______ zeff values pull electrons ____ strongly toward the nucleus. Therefore, _____ energy is required to rip them off

Moving down a group, Zeff _____, but the distance of the electron from the nucleus _____ as well. Coulomb's law demonstrates that the electrostatic force, F, ______ with the square of the distance from the nucleus, r. Due to the exponent, the increased distance plays a ____ important role than the increased Zeff and the attractive electrostatic force ______ down a group. Less force means _____ energy is required to remove the electron, so ionization energy ______ moving down a group.

Define electronegativity
When two atoms have different electronegativities, they share ______ unequally, causing ______. Relative electronegativity determines the direction of ______ within a bond and within an overall molecule.

Like ionization energy, electronegativity tends to increase across a period from ____ to _____ and ____ a group. The most commonly used measurement of electronegativity is the ______ scale, which ranges from a value of 0.79 for cesium to a value of 4.0 for fluorine. It is important to remember that _____ is the most electronegative element.

The electronegativity of hydrogen falls between that of _____ and that of _____. When bonded with hydrogen, carbon and elements to the right of carbon will carry a partial ______ charge while hydrogen will carry a partial ______ charge. Think of CH4. Boron and the elements to the left of boron will carry a partial ______ charge when bonded to hydrogen, while the hydrogen will carry a partial ______ charge. Think of the hydrides. (H^-) in NaH or LiAIH3

Since noble gases tend not to make bonds, electronegativity values are ______ for the noble gases. Electronegativity values provide a system for predicting which type of bond will form between two atoms. Atoms with large differences in electronegativity (__ or larger on the Pauling scale as a rule of thumb) will form _____ bonds. Metals and non-metals usually exhibit ____ electronegativity differences and form ____ bonds with each other.

Atoms with moderate differences in electronegativities (___ -___ on the Pauling scale) will generally form _____ _____ ____. Atoms with very minor electronegativity differences (___ or smaller on the Pauling scale) will form ______ _____ bonds.

Define Electron affinity

Just like electronegativity, electron affinity tends to increase on the periodic table from _____ to _____ and from ______ to ____. The sign of electron affinity values can be different for different atoms (explain)

Warning: Electron affinity is sometimes described in terms of exothermicity, for which the energy released is given a negative sign. We can state this as follows: electron affinity is more exothermic to the right and up on the periodic table. The noble gases ____ follow this trend. Electron affinity values for the noble gases are ______ (why)

Neils Bohr proposed a theory of the atom, known as the Bohr atom, which represents the atom as a nucleus surrounded by electrons in discrete _____ _____. In the orbital structure of the hydrogen atom, a single electron orbits the hydrogen nucleus in an _____ _____. Any electron in any atom of any element, including this hydrogen electron, can be described by four ______ ______, which serve as that electron's "ID number." State the Pauli Exclusion Principle

The first quantum number is the ____ _____ _____ (__). (Define)

The second quantum number is the _______ ______ _____ (__). (Define)

The third quantum number is the _____ _____ ______ (__). (Define)

Within the p-subshell (l = 1), there are three magnetic quantum numbers: ___, ___, and ___. An atom will spread out its electrons amongst the orbitals, such that all orbitals in a subshell have ___ electron before any has ___. The number of total orbitals within a shell is equal to ___. Solving for the number of orbitals for each shell gives 1, 4, 9, 16 ... Since there are ____ electrons in each orbital, the number of elements in the periods of the periodic table is ___, ___, ___, and ___.

The fourth quantum number is the _____ _____ _____ _____ (__). The electron spin quantum number has possible values of___ or ___. This final quantum number is used to distinguish between the ___ electrons that may occupy the same orbital, since those electrons will have the same first ____ quantum numbers.

The size of an atom has a significant effect on its chemistry. Small atoms hold charge in a concentrated way because they have _____ orbitals available to distribute and thereby ______ charge. This concentration of charge makes the ______ element in each group bonds _____ readily and with _____ bond strength, especially when in _____ form. Fluorine is a good example. The fluoride ion (F) is too _____ to manage its full negative charge. Therefore, it is generally insoluble and bonds immediately when in solution. For this reason, in toothpaste, fluoride (a poison in high concentrations) _____ immediately with the enamel of teeth before it can be ingested.

Small atoms do not have d-orbitals available for bond formation, and therefore cannot form more than ___ bonds. Large atoms have d-orbitals, allowing for more than __ bonds. Oxygen typically forms two bonds, while the Larger sulfur can form up to ___. Smaller atoms have the advantage when it comes to __-bonding. The p-orbitals on large atoms do not _____ significantly, so they cannot easily form __-bonds. Carbon, nitrogen, and oxygen are small enough to form ____ π-bonds while their Larger third row family members form only ____ π-bonds, if they form π-bonds at all.

Let's summarize: The flrst quantum number is the ____. It corresponds roughly to the energy level of the _____ within that shell. The second quantum number is the ______. It gives the _____. Recognize that s orbitals are ______ and p orbitals are ______-shaped. The third quantum number'is the specific _____ within a subshell. The fourth quantum number distinguishes between two ______ in the same orbital; one is ____ ____ and the other is ____ ____.

State the Name, symbol, value and character of the quantum numbers

State Heisenberg's Uncertainty principle
This uncertainty arises from the dual nature (wave-particle) of matter, and is on the order of Planck's constant (6.6 x 10-34 J s):

To put the uncertainty principle more simply, the more we know about the _______ of any particle, the less we can know about its ______, and vice versa. There are other quantities besides position and momentum to which the uncertainty principle applies, but position and momentum are the ones that will likely be tested on the MCAT®.

Define the Aufbau principle

All other things being equal, the lower the energy state of a _____, the more stable the system. Thus, electrons look for an orbital in the _____ energy level whenever they are added to an atom. The orbital with the ______ energy will be located in the ______ with the lowest energy.

For the representative elements, the shell level of the most recently added electrons is given by the _____ in the periodic table. The most recently added electron for Sr is in shell __, and the most recently added electron for Sr is in shell __.

For transition metals the shell of the most recently added electron lags ____ behind the period. The most recently added electron of Ag is in shell 4, though it is in the ____ period. For the lanthanides and actinides, the shell of the most recently added electron lags ____ behind the period. The most recently added electron for Ce is in shell 4, though the lanthanide series corresponds to the ____ period.

The subshells (s, p, d, and}) are the orbital shapes with which we are familiar. Orbital shapes are not true paths that  electrons follow, but rather represent _____ _____ for the position of an _____. There is a 90% chance of finding the electron somewhere inside the given shape. Be familiar with the shapes of the orbitals in the __- and __-subshells. The shapes of the d- and f-subshells are beyond the scope of the MCAT®

The organization of the periodic table indicates where each type of _____ is filling. The s-subshells are filling in groups ___ and ___, the p-subshells are filling in groups ___-___, the d-subshells are filling in groups ___-___, and the f-subshells are filling in the _______ and ______ series. Therefore, the periodic table can be used to determine the _____ and _____ of the most recently added electron for any element. For example, the most recently added ______ in S is in the 3p energy level; for Au it is in the ___ energy level; and forK it is in the ___ energy level.

Valence electrons, the electrons which contribute most to an element's ______ properties, are located in the ______ _____ of an atom. In most cases, only electrons from the ___ and ___ subshells are considered valence electrons.