Chemistry Unit II Study Guide

Understand the contributions of J. J. Thomson, Robert Millikan, and Ernest Rutherford to atomic theory

1. In 1897, English physicist J. J. Thomson experimented with a device called a cathode ray tube, in which an electric current was passed through gases at low pressure. Thomson conducted further experiments which allowed him to calculate the charge-to-mass ratio (e/me) of the electron. He found that this value was a constant and did not depend on the gas used in the cathode ray tube or on the metal used as the electrodes. He concluded that electrons were negatively charged subatomic particles present in atoms of all elements. 2. American physicist, Robert Millikan, carried out a series of experiments between 1908 and 1917 that allowed him to determine the charge of a single electron. Millikan's experiment was called the oil drop experiment Millikan's oil drop experiment: Oil drops that are sprayed into the main chamber fall through a tiny hole into an electric field, after which they can be viewed through a microscope. This experiment allowed Millikan to determine the charge of the electron. he was also able to calculate the mass of a single electron. Charge of one electron=−1.602×10−19 C Mass of one electron=9.11×10−28 g 3. In 1911, Rutherford and coworkers Hans Geiger and Ernest Marsden initiated a series of groundbreaking experiments that would completely change the accepted model of the atom. The experimental setup for Rutherford's gold foil experiment: A radioactive element that emitted alpha particles was directed toward a thin sheet of gold foil, which was surrounded by a screen that would allow detection of the deflected particles. According to the plum pudding model all of the alpha particles should have passed through the gold foil with little or no deflection. Rutherford found that a small percentage of alpha particles were deflected at large angles, which could be better explained by an atom that contained a very small, dense, positively-charged nucleus He concluded that all of the positive charge and the majority of the mass of the atom must be concentrated in a very small space in the atom's interior, which he called the nucleus. Rutherford's atomic model became known as the nuclear model. However, it did not completely address the nature of the electrons and the way in which they occupied the vast space around the nucleus.

Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.

1. law of conservation of mass, which states that during a chemical reaction, the total mass of the products must be equal to the total mass of the reactants. In other words, mass cannot be created or destroyed during a chemical reaction, but it must always be conserved. 2. law of definite proportions, which states that a given chemical compound always contains the same elements in the exact same proportions by mass. 3. law of multiple proportions. Whenever the same two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.

Describe some of the early attempts to organize the chemical elements.

A logical way to begin to group elements together was by their chemical properties. In other words, putting elements in separate groups based on how they reacted with other elements. In 1829, a German chemist named Johann Dobereiner (1780-1849) placed various groups of three elements into groups called triads. Triads were based on both physical and chemical properties. English chemist John Newlands (1838-1898) arranged the elements in increasing order of atomic mass and noticed that every eighth element exhibited similar properties. He called this relationship the Law of Octaves.

Describe various components of the modern periodic table, including periods, groups, metals, nonmetals, and metalloids.

A period is a horizontal row of the periodic table. A group is a vertical column of the periodic table. A metal is an element that is a good conductor of heat and electricity. A nonmetal is an element that is generally a poor conductor of heat and electricity.

Describe how changes in electron energies lead to atomic emission spectra.

An atomic emission spectrum is the pattern of lines formed when the light emitted from an atom passes through a prism to separate it into the different frequencies of light it contains. {An atom is in an excited state when its potential energy is higher than that of the ground state. An atom in an excited state is not stable. When it returns back to the ground state, it releases the energy that it had previously gained in the form of electromagnetic radiation.}

Understand how Mendeleev organized his periodic table.

He famously organized the information for each element onto separate note cards that were then easy to rearrange as needed. He discovered that when he placed them in order of increasing atomic mass, certain similarities in chemical behavior repeated at regular intervals. This type of a repeating pattern can be referred to as periodic.

Be able to calculate the average atomic mass of an element.

An atomic mass unit is defined as a mass equal to one twelfth the mass of an atom of carbon-12. Known chlorine-35: atomic mass = 34.969 amu and % abundance = 75.77% chlorine-37: atomic mass = 36.966 amu and % abundance = 24.23% Unknown Average atomic mass of chlorine Change each percent abundance into decimal form by dividing by 100. Multiply this value by the atomic mass of that isotope. Add together the results for each isotope to get the average atomic mass. Step 2: Calculate. chlorine-35chlorine-37average atomic mass0.7577×34.969=26.50 amu0.2423×36.966=8.957 amu26.50+8.957=35.45 amu Note: Applying significant figure rules results in the 35.45 amu result without excessive rounding error. In one step: (0.7577×34.969)+(0.2423×36.966)=35.45 amu Step 3: Think about your result. The calculated average atomic mass is closer to 35 than to 37 because a greater percentage of naturally occurring chlorine atoms have a mass number of 35. It agrees with the value listed in the table above

Describe how ions are formed.

An ion is an atom or group of bonded atoms that has a positive or negative charge. How do atoms obtain this charge? In theory, there are two possibilities: (1) gaining or losing positively charged protons, or (2) gaining or losing negatively charged electrons. However, the nucleus of an atom is very stable, and the number of protons cannot be changed by chemical reactions.

Learn the periodic trends for electronegativity.

Electronegativity is a measure of the ability of an atom to attract shared electrons when the atom is part of a compound. Electronegativities generally increase from left to right across a period. This is due to an increase in nuclear charge. Alkali metals have the lowest electronegativities, while halogens have the highest. Because most noble gases do not form compounds, they are generally not assigned electronegativity values. Note that there is little variation among the transition metals. Electronegativities generally decrease from top to bottom within a group due to the larger atomic size.

Know the improvements that Moseley made to Mendeleev's table.

English physicist Henry Moseley examined the x-ray spectra of a number of chemical elements. His results led to the definition of atomic number as the number of protons contained in the nucleus of each atom. He then realized that the elements of the periodic table should be arranged in order of increasing atomic number instead of increasing atomic mass.

Describe the Bohr model of the atom.

In 1913, Danish physicist Neils Bohr (1885-1962) proposed a model of the atom that explained the hydrogen atomic emission spectrum. According to the Bohr model, which is often referred to as a planetary model, the electrons encircle the nucleus of the atom in specific allowable paths called orbits. When an electron is in one of these orbits, its energy is fixed. The ground state of the hydrogen atom, where its energy is lowest, occurs when its single electron is in the orbit that is closest to the nucleus. The orbits that are farther from the nucleus are all higher energy states. The electron is not allowed to occupy any of the spaces in between the orbits.

Understand the de Broglie wave equation and how it illustrates the wave nature of the electron.

In 1924, French scientist Louis de Broglie (1892-1987) derived an equation that described the wave nature of any particle. He determined that the wavelength (λ) of any moving object is given by: λ=hmv In this equation, h is Planck's constant, m is the mass of the particle in kg, and v is the velocity of the particle in m/s. The problem below shows how to calculate the wavelength of an electron.

Predict the effect that ion formation has on the size of an atom.

In most cases, the formation of an anion by the addition of an electron to a neutral atom releases energy. The addition of electrons always results in an anion that is larger than the parent atom.

Describe the structure of the nuclear atom.

In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom. How can you describe the structure of the nuclear atom?

Learn the periodic trends for ionization energy.

Ionization energy is the energy required to remove an electron from an atom. An equation can be written to illustrate this process for a sodium atom: Na+energy→Na++e− The ionization energies of various elements are influenced by the size of the atom, the nuclear charge, and the electron energy levels. Ionization energies are measured in units of kilojoules per mole (kJ/mol).

Understand how isotopes differ from one another and be able to designate them by various methods.

Isotopes - atoms that have the same atomic number but different mass numbers due to a change in the number of neutrons.

Describe John Dalton's atomic theory.

John Dalton (1766-1844) formulated an atomic theory based on the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. His theory can be summarized in the following statements, 1. All matter is composed of extremely small particles called atoms. 2. Atoms of the same element are identical in terms of size, mass, and other properties. Atoms of one element are different from the atoms of any other element. 3. Atoms of different elements can chemically combine with one another in simple whole-number ratios to form chemical compounds. 4. Chemical reactions can be described as the separation, combination, or rearrangement of atoms. Atoms of one element cannot be changed into atoms of a different element as a result of a chemical reaction.

Explain the difference between quantum mechanics and classical mechanics.

Quantum mechanics is the study of the motion of objects that are atomic or subatomic in size and thus demonstrate wave-particle duality. In classical mechanics, the size and mass of the objects involved effectively obscures any quantum effects, so such objects appear to be capable of gaining or losing energy in any amount. Particles whose motion is better described by quantum mechanics can only gain or lose energy in discrete units called quanta. (Another feature that is unique to quantum mechanics is the uncertainty principle.)

Understand that some electron configurations are exceptions to the normal Aufbau process.

Some electron configurations, such as those of chromium and copper, do not strictly follow the Aufbau principle.

Understand how the Heisenberg uncertainty principle and Schrödinger's wave equation led to the idea of atomic orbitals.

The Heisenberg Uncertainty Principle states that it is impossible to determine simultaneously both the position and the velocity of a particle. In 1926, Austrian physicist Erwin Schrödinger (1887-1961) used the wave-particle duality of the electron to develop and solve a complex mathematical equation that accurately described the behavior of the electron in a hydrogen atom. The Heisenberg uncertainty principle led to the idea of atomic orbitals because it stated that it was impossible to simultaneously determine both the position and the velocity of an electron or any other particle, allowing the idea of an electron following an orbital path where a general area could be determined. The Schrodinger wave equation treated electrons as waves and had an outcome of quantization of electron energies. Together, both the principle and the wave equation laid the foundation for modern quantum theory.

Define atomic number.

The atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.

Learn the periodic trends for atomic radius.

The atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together Atomic radius generally decreases from left to right across a period, although there are some small exceptions to this trend, such as the relative radii of oxygen and nitrogen.

Know how to use the noble gas notation shorthand method.

The electron configuration of an atom can be abbreviated by using noble gas notation, in which the elemental symbol of the last noble gas prior to that atom is written first, followed by the configuration of the remaining electrons. Electron Configurations of Third-Period Elements Element Name Symbol Atomic Number Electron Configuration Sodium Na 11 [Ne]3s1 Magnesium Mg 12 [Ne]3s2 Aluminum Al 13 [Ne]3s23p1 Silicon Si 14 [Ne]3s23p2 Phosphorus P 15 [Ne]3s23p3 Sulfur S 16 [Ne]3s23p4 Chlorine Cl 17 [Ne]3s23p5 Argon Ar 18 [Ne]3s23p6 *the number of valence electrons increases from one to eight across the third period.

Locate the following groups on the periodic table: alkali metals, alkaline earth metals, halogens, and noble gases.

The elements in Group 1 (lithium, sodium, potassium, rubidium, cesium, and francium) are called the alkali metals. The elements in Group 2 (beryllium, magnesium, calcium, strontium, barium, and radium) are called the alkaline earth metals The elements of Group 17 (fluorine, chlorine, bromine, iodine, and astatine) are called the halogens The elements of Group 18 (helium, neon, argon, krypton, xenon, and radon) are called the noble gases

Describe electron affinity.

The energy change that occurs when a neutral atom gains an electron is called its electron affinity.

Explain how multiple ionization energies are related to noble gas electron configurations.

The first ionization energies for the noble gases (He, Ne, Ar) are higher than those of any other element within that period. The noble gases have full outer s and p sublevels, which gives them extra stability and makes them mostly unreactive.

Describe the relationship between outer electron configuration and group number. Be able to determine the number of valence electrons for any element.

The group number of the representative elements = the number of valence electrons. The representative elements are in columns A1 - A7. In an 18 column table the map is as follows: A1 = column 1, hydrogen + alkali metals have 1 valence A2 = column 2, alkaline earth metals have 2 valence A3 = column 13, Boron family have 3 valence A4 = column 14, Carbon group have 4 valence A5 = column 15, Nitrogen family have 5 valence A6 = column 16, Oxygen column - chalcogens have 6 valence A7 = column 17, Fluorine column - halogens have 7 valence

Understand the photoelectric effect and how it is related to the wave-particle duality of light.

The photoelectric effect is a phenomenon that occurs when light shined onto a metal surface causes the ejection of electrons from that metal. wave-particle duality of light - proposed that light could also be described as quanta of energy that behave as particles. Einstein used the particle theory of light to explain the photoelectric effect

Understand how to apply the Aufbau principle to determine ground state electron configurations.

To determine the ground state electron configuration for a given atom, it is first necessary to organize the atomic sublevels in order of increasing energy. The Aufbau principle states that all lower energy orbitals must be filled before electrons can be added to a higher energy orbital. The Aufbau principle is sometimes referred to as the "building-up" principle.

Locate the transition elements, lanthanides, and actinides on the periodic table.

Transition elements are the elements that are found in Groups 3-12 on the periodic table. The lanthanides are the 14 elements from cerium (atomic number 58) to lutetium (atomic number 71). The actinides are the 14 elements from thorium (atomic number 90) to lawrencium (atomic number 103).

Know the relationship between group number and valence electrons.

Valence Electrons for the Representative Elements Group Number Outer Electron Configuration Number of Valence Electrons 1 ns1 1 2 ns2 2 13 ns2np1 3 14 ns2np2 4 15 ns2np3 5 16 ns2np4 6 17 ns2np5 7 18 ns2np6 8 You can see that the number of valence electrons and the outer electron configuration is constant within a group. This is the reason why elements within a group share similar chemical properties.

Be able to determine the number of valence electrons and the number of unpaired electrons in any atom.

Valence electrons are the electrons in the outermost principal energy level. An atom can have a maximum of eight valence electrons. ?

Describe the relationships between speed, wavelength, and frequency of light.

Wavelength and frequency are therefore inversely related. As the wavelength of a wave increases, its frequency decreases. The equation that relates the two is: c=λν The variable c is the speed at which the wave is traveling. (In the case of electromagnetic radiation, c is equal to the speed of light.)

Be able to write correct electron configurations for all elements.

Writing the electron configuration of an atom essentially amounts to listing which orbitals contain electrons and how many electrons are in each type of orbital.

Define mass number.

mass number is defined as the total number of protons and neutrons in an atom.

Know the periodic law.

periodic law, which states that when elements are arranged in order of increasing atomic number, there is a periodic repetition of their chemical and physical properties.

Know the four quantum numbers and how they are related to the arrangement of electrons in an atom.

principle quantum number, angular momentum qn, magnetic qn, spin qn The principal quantum number is symbolized by the letter n and designates the principal or main energy level occupied by the electron. The angular momentum quantum number is symbolized by the letter l and indicates the shape of the orbital. The magnetic quantum number is symbolized by the letter ml and indicates the orientation of the orbital around the nucleus. The spin quantum number is symbolized by the letter ms and indicates the direction of electron spin.

Distinguish between the three main subatomic particles.

protons, electrons, and neutrons

Identify each block of the periodic table and be able to determine which block each element belongs to based on its electron configuration.

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Understand the relationship between the number of orbitals in various energy sublevels and the length of the periods in the periodic table.

the four different sublevels (s, p, d, and f) each consist of a different number of orbitals. The s sublevel has one orbital, the p sublevel has three orbitals, the d sublevel has five orbitals, and the f sublevel has seven orbitals. In the first period, only the 1s sublevel is being filled. Since all orbitals can hold two electrons, the entire first period consists of just two elements. In the second period, the 2s sublevel, with two electrons, and the 2p sublevel, with six electrons, are being filled. Consequently, the second period contains eight elements. The third period is similar to the second, except the 3s and 3p sublevels are being filled. Because the 3d sublevel does not fill until after the 4s sublevel, the fourth period contains 18 elements, due to the 10 additional electrons that can be accommodated by the 3d orbitals. The fifth period is similar to the fourth. After the 6s sublevel fills, the 4f sublevel is populated with up to 14 electrons. This is followed by the 5d and the 6p sublevels. The total number of elements in the sixth period is 32. The seventh period also contains 32 elements, most of which are too unstable to be found in nature. All 32 have been detected or synthesized, although, for some of the later elements in this period, only a handful of atoms have ever been made.