Ultimate chemical equations handbook answers chapter 12

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Thomson concluded that the ray consisted of a stream of negatively charged particles that had been dislodged from atoms. These particles became known as electrons. Thomson went on to determine that the mass of an electron is much less than that of a hydrogen atom. This value was accepted as a fundamental unit of electrical charge. In , Ernest Rutherford conducted experiments in which he aimed a beam of alpha particles positively charged helium nuclei at a thin sheet of gold.

This positively charged mass is called the nucleus of the atom. In , Rutherford and Chadwick demonstrated the existence of the neutron, a nuclear particle having no charge but with nearly the same mass as a proton. The following table summarizes the sub- atomic particles that comprise the atom. Select the term in Column B that best matches the phrase in Column A.

Column A Column B a. Cathode ray 1. Discovered in 2. Caused large deflections of alpha 3. Has no charge f. In an atom, the number of these particles is equal to the number of protons.

Identified by Thomson i. Has a positive charge and relative mass of 1 k. The center of an atom l. Starting with the simplest element, hydrogen, the periodic table represents each element in a box.

As you scan the table, you can see that the elements are arranged in order of increasing atomic number and that the sequence of atomic numbers has no missing values.

This continuous series is an indication of the fact that the atomic number determines the identity of an element. Thus, positive charges protons and negative charges electrons balance out. Electrons, protons, and atomic number are related as follows.

How many electrons and protons are contained in an atom of each of the following elements? Identify the atom containing the following number of electrons. Identify the atom containing the following number of protons. Atoms of which element contain 18 protons and 18 electrons?

For example, because the atomic number of hydrogen is one, all atoms of hydrogen must have one proton and one electron. However, hydrogen atoms can have different numbers of neutrons. Approximately 99 out of every naturally occurring hydrogen atoms have no neutrons. However, the remaining 15 out of every hydrogen atoms have one neutron. Atoms with the same number of protons but dif- ferent numbers of neutrons are called isotopes.

Hydrogen has two naturally occurring isotopes. H atom with one proton H atom with one proton and no electrons and one neutron Mass number 1 Mass number 2 Atomic number 1 H Atomic number 1 H Isotopes of an element can also be specified by writing the mass number of the isotope following a hyphen after the name of the element.

Substituting atomic number for the number of pro- tons gives the following relationship. You can use this relationship to determine any one of the variables if you know the other two. Example Problem Using Atomic Number and Mass Number One of the four naturally occurring isotopes of chromium has a mass number of Determine the number of protons, electrons, and neu- trons in an atom of this isotope and write its symbol.

Start by obtaining information on chromium from the periodic table. Chromium Cr is found in the fourth row of the table in group 6. The atomic number of chromium is You know that the atomic number of an element gives the number of protons in the nucleus of the atom. Mass number a Atomic number b X Symbol of element Substituting the known values gives this symbol for the isotope.

The other three naturally occurring isotopes of chromium have mass numbers of 50, 52, and Problem All naturally occurring atoms of fluorine have a mass number of Describe the subatomic particles comprising an isotope of zirconium 94 40 Zr. An atom of a certain element has a mass number of and is known to contain 64 neutrons.

Identify the atom and determine the number of electrons and protons the atom contains. A neutral atom has 78 electrons and a mass number of Identify the atom and determine the number of protons and neutrons in its nucleus.

The number at the bottom of each square in the periodic table is the atomic mass of that element in amu. A scan of the periodic table will immediately tell you that the atomic masses of many elements are not whole numbers. The reason for this is that many elements occur in nature as a mixture of isotopes—isotopes with different masses. The atomic mass of an element given in the table is a weighted average of the atomic masses of the naturally occurring isotopes of the element.

This weighted average takes into account the mass and abundance of each of the isotopes. The following Example Problem shows you how to calculate the atomic mass of an element. The lighter isotope Cu , with 29 protons and 34 neutrons, makes up The atomic mass of Cu is Use the data above to compute the atomic mass of copper.

For a problem that has several numerical values, it is wise to organize the data into a table. Gallium occurs in nature as a mixture of two isotopes. Ga with a Calculate the atomic mass of gallium. The following table shows the five isotopes of germanium found in nature, the abundance of each isotope, and the atomic mass of each isotope.

Calculate the atomic mass of germanium. The atomic mass of bromine given in the periodic table is Use a reference book to find the percent of Br in naturally occurring bromine. Explain the value of the atomic mass of bromine from the data you find. The element chlorine occurs in nature as a mixture of two iso- topes.

Chlorine has an atomic mass of Chlorine atoms make up the remaining Use the average atomic mass of chlorine from the periodic table to calculate the atomic mass of Cl atoms. However, scientists in the late nineteenth century began to notice that some elements spon- taneously emitted energy and particles.

These emissions became known as radiation. Elements that give off radiation are said to be radioactive. Soon, scientists discovered that radioactive elements changed into other elements after emitting radiation. Thus, atoms were not unchangeable as Dalton had thought. Today, we know that some nuclei are unstable and gain stability by emitting radiation. This process is called radioactive decay. One type is alpha radiation -radiation , which consists of alpha particles ejected from the nucleus.

An alpha particle is equivalent to a helium-4 nucleus, that is, a nucleus with two protons and two neutrons. Gamma radiation is usually emitted along with -particles and -particles. Chapter 4 Review Why was the discovery of the electron considered evidence for the existence of positively charged particles in the atom? What particles are responsible for that mass? Explain the relationship between the number of protons and the number of electrons in an atom of argon.

Can atoms of two different elements have the same number of protons? What particles make up an atom of selenium? Explain the meaning of the number printed at the bottom of each box of the periodic table. Write symbols for the following isotopes. Complete the following table. X rays, gamma rays, and radio waves are other examples of electromagnetic radiation. Electromagnetic waves, along with water waves and sound waves, exhibit certain common characteristics.

All waves consist of a series of crests and troughs that travel away from their source at a velocity that is determined by the nature of the wave and the mate- rial through which the wave passes.

The frequency and velocity of a wave determine its wavelength, the distance between equivalent points on a continuous wave. For waves of a given velocity, a wave with a higher frequency has a shorter wavelength.

For electromagnetic waves, this inverse relation is expressed mathematically by the following equation. A helium-neon laser emits light with a wavelength of nm. What is the frequency of this light? What is the wavelength of X rays having a frequency of 4. An FM radio station broadcasts at a frequency of Whether or not electrons are ejected depends on the frequency color of the incident light.

In , Einstein reasoned that this phe- nomenon could be explained only if light could act like a stream of particles that knocked electrons out of atoms. Particles, or photons, of light at the high-frequency violet end of the visible spectrum had greater energy and were therefore more effective at dislodging electrons.

Practice Problems 4. Calculate the energy of a gamma ray photon whose frequency is 5. What is the difference in energy between a photon of violet light with a frequency of 6. Calculate the energy of a photon of ultraviolet light that has a wavelength of A neon sign works on this principle. This emitted light can be broken into a spectrum consisting of discrete lines of specific frequencies, or colors. Niels Bohr used the relatively simple emission spectrum of the hydrogen atom to propose a new atomic model.

When the electron is in the orbit nearest the nucleus, it has the lowest possible energy. When the electron is in the next larger orbit, it has a higher energy, and so on through the larger allowed orbits. The electron can occupy only those specific allowed orbits and thus can have only the energies associated with those orbits.

When it falls back to a lower-energy orbit, the electron emits an amount quantum of energy equal to the differ- ence in energy between the two orbits. Because the emission spectrum of hydrogen consisted of several different frequencies, Bohr assumed that the electron could occupy orbits of several differ- ent energies, designated by integers called quantum numbers. However, the model was very limited in that it only worked for hydrogen.

In , Louis de Broglie turned the tables by theorizing that particles of matter, specifically electrons, could exhibit the prop- erties of waves—frequency and wavelength. De Broglie proposed that electrons of only a specific frequency could fit into one of the possible atomic orbits.

This model may be summarized as follows. Electrons occupy the space surrounding the nucleus and can exist in several discrete princi- pal energy levels, each designated by one of the principal quantum numbers n that are the integers 1, 2, 3, 4, and so on.

Electrons in successively higher principal energy levels have greater energy. Because of interactions among electrons, each principal energy level consists of energy sub- levels that have slightly different energy values.

These sublevels are designated by the letters s, p, d, and f in order of increasing energy. Each energy sublevel consists of one or more orbitals, each of which can contain two electrons. An s sublevel has one orbital; a p sublevel has three orbitals; a d sublevel has five orbitals; and an f sublevel has seven orbitals. All of the orbitals in the same sublevel are of equal energy.

Atomic orbitals are regions of space in which there is a high probability 90 percent of finding an electron. Except for s orbitals, the orbitals are not spherical in shape. Instead, they can be anywhere within the orbital, and there is a 10 percent chance that the electron will be located outside the orbital.

Describe what is happening when an atom emits a photon. How many electrons can the second principal energy level hold? How many electrons can the third principal energy level hold? Explain the difference in these numbers of electrons. Therefore, it is useful to be able to determine and write out the elec- tron arrangement, called an electron configuration, of an atom in symbolic form. The number of electrons in an energy sublevel is indicated by a super- script integer.

Helium has two electrons. Recall that an s sublevel consists of a single orbital that can hold a maximum of two electrons. A neutral lithium atom has three electrons. The first principal energy level is filled with two electrons. Where does the third elec- tron go? This electron is found in the second principal energy level, which, like all energy levels, begins with an s sublevel. So, lithium has two electrons in the 1s orbital and a third electron in the 2s orbital, giving it the electron configuration 1s22s1.

With beryllium, which has four electrons, the 2s orbital is filled with two electrons, yielding a configuration of 1s22s2. Recall that the second principal energy level has two energy sublevels available s and p and that the p sublevel is of higher energy than the s sublevel.

One rule gov- erning electron configurations is the aufbau principle, which states that each successive electron occupies the lowest energy orbital available. In the case of boron, the lower-energy s orbitals are full, so the fifth electron is found in one of the three available orbitals in the higher-energy 2p sublevel.

Recall that the three orbitals available in a p sublevel can each hold a maximum of two electrons. Continuing from boron through carbon 1s22s22p2 , nitrogen 1s22s22p3 , oxygen 1s22s22p4 , fluorine 1s22s22p5 , and neon 1s22s22p6 , the 2p sublevel becomes filled with six electrons. With the element sodium, the eleventh electron begins the 3s sublevel to give the configuration 1s22s22p63s1.

The same pattern that occurred with lithium through neon repeats here, with succes- sive electrons filling the 3s orbital and 3p orbitals. The element argon has a filled 3p sublevel. Argon has 18 total electrons, and the configuration 1s22s22p63s23p6.

You may recall that three sublevels—s, p, and d—are available in the third energy level, so you might expect the next electron to begin the 3d sublevel. However, a complication occurs here because the 4s sublevel is of lower energy than the 3d sublevel. Thus, fol- lowing the aufbau principle, the next nineteenth electron begins the 4s sublevel in the element potassium, producing the configuration 1s22s22p63s23p64s1. Scandium is the first element that has electrons in the 3d sublevel.

As an example of an element that has electrons in the 4p sublevel, the electron configuration of arsenic is 1s22s22p63s23p64s23dp3. Notice that the configuration is written by placing the sublevels in order of increasing energy, not in numerical order. When writing the electron configuration of an atom, in what general order are the sublevels written?

How is the number of electrons in an energy sublevel indicated in an electron configuration? Write the electron configurations of the following elements.

By following the arrows, you can write the energy sublevels in the correct order. Helium, the preceding noble gas, has the configuration 1s2.

This form of an electron configuration is called noble-gas notation. Example Problem Writing Electron Configurations Using Noble-Gas Notation Use noble-gas notation to write the electron configuration for a neu- tral atom of silicon.

To begin, find silicon on the periodic table and locate the preceding noble gas. Silicon Si is atomic number 14, and the preceding noble gas is neon Ne , atomic number Start the electron configuration by writing the symbol [Ne]. Next, note that silicon has four more electrons than neon. By combining this information, you can write the following electron configuration for silicon. Using noble-gas notation, write the electron configurations of the following elements.

Therefore, these outermost electrons, called valence elec- trons, determine most of the chemical properties of an element. The American chemist G. In writing electron-dot structures, a single dot is used to repre- sent each valence electron. One dot is placed on each of the four sides around the symbol before any two dots are paired together. The following examples illustrate the process. The electron-dot structure for hydrogen, which has one electron, is H.

The electron- dot structure for helium, 1s2, is He. The electron configuration of lithium is 1s22s1, but the two 1s electrons are in a stable inner energy level and do not participate in chemical changes. Only the outermost 2s electron is a valence electron, so the electron-dot structure for lithium is Li. The electron-dot struc- ture for boron, with three electrons in the second energy level 1s22s22p1 , is B.

The electron configuration of oxygen is 1s22s22p4, and its electron-dot structure is O. A new principal energy level begins with sodium, whose electron configuration is 1s22s22p63s1. The other electrons are in inner energy levels. Example Problem Writing Electron-Dot Structures Write an electron-dot structure for a neutral atom of arsenic, an element used in semiconductor materials because of its electron con- figuration.

Begin by writing the electron configuration of arsenic. Use noble- gas notation to emphasize the arrangement of electrons in the highest principal energy level. Arsenic has 15 electrons more than argon. Using the energy sublevel diagram, you can see that the nineteenth electron begins the fourth principal energy level. After the 4s sub- level is filled with two electrons, ten electrons fill the 3d sublevel, which is slightly higher in energy than the 4s sublevel.

After the 3d sublevel is filled, the remaining three electrons will be in orbitals in the 4p sublevel. Thus, the electron configuration for arsenic is [Ar]4s23dp3. Therefore, they are not regarded as valence electrons.

Only the 4s and 4p electrons are valence electrons; thus, the electron-dot structure of arsenic has five dots as follows. Write electron-dot structures for the following elements. What electron-dot structure is shared by all noble gases except helium? List three elements that have the electron-dot structure X. Chapter 5 Review When light passes from air into a denser material, such as glass or water, it slows down. Describe the photoelectric effect.

How does the energy of a photon of electromagnetic energy change as the frequency increases? As the wavelength increases?

How many energy sublevels are available in the third principal energy level? How many electrons can each of these sublevels hold? Identify the elements that have the following electron configurations. What electrons do the dots in an electron-dot structure represent? Why are these electrons important? Mendeleev made a table in which he arranged the elements in order of increasing atomic mass into columns with similar properties.

From his table, Mendeleev predicted the properties of three elements that were undiscovered at the time. Each square gives certain information about each element, as shown in the following diagram. Because the pattern of properties repeats in each new row of elements, the ele- ments in a column have similar properties and are called a group or family of elements. The groups are designated with a number and the letter A or B.

Groups 1A through 8A are called the main group or representative elements. The group B elements are called the transition elements. Elements are divided into three main classes—metals, metal- loids, and nonmetals. As you can see from the periodic table, the majority of the elements are metals. Metals are generally shiny solids and are good conductors of heat and electrical current. Some groups of elements have names.

For example, the first two groups of metals, groups 1A and 2A, are called the alkali metals and the alkaline earth metals, respectively. Most of the elements to the right of the heavy stair-step line in the periodic table are nonmetals, which are generally either gases or brittle solids at room tempera- ture. Group 7A elements are commonly called halogens, and group 8A elements are the noble gases.

Many of the elements that border the stair-step line are metalloids, which have some of the character- istics of both metals and nonmetals. Each Column A element may match more than one description from Column B. Column A Column B 1.

Thus, the arrangement of elements in the periodic table reflects the electron structures of atoms. For example, the group number of a representative element gives the number of valence electrons in an atom of that element. The group 3A element aluminum, for example, has the electron configu- ration 1s22s22p63s23p1, or [Ne]3s23p1. The three electrons in the third energy level 3s2 and 3p1 are the valence electrons of the alu- minum atom.

In a similar way, the period number of a representative element indicates the energy level of the valence electrons. Practice Problems Use the periodic table to answer the following questions. How many valence electrons are in an atom of each of the following elements? In which energy level are the valence electrons of the elements listed in question 8?

Identify each of the following elements. The 1A and 2A groups constitute the s-block elements because their highest-energy electrons are in s orbitals. The remaining groups of representative elements, 3A through 8A, make up the p-block of elements.

In these elements, s orbitals are filled, and the highest-energy electrons are in p orbitals. The transition metals are the d-block elements. In these ele- ments, the highest-energy electrons are in the d sublevel of the energy level one less than the period number. The remaining block is the f-block, or inner transition metals. The highest-energy electrons in these elements are in an f sublevel of the energy level two less than the period number.

The electron configuration of phosphorus is [Ne]3s23p3. Without using the periodic table, determine the group, period, and block in which the element is located in the periodic table. First, identify the valence electrons and note their energy level. In phosphorus, the 3 in front of the s and p orbitals indicates that the valence electrons are in the third energy level.

Therefore, phospho- rus will be found in the third period of the periodic table. Next, note the sublevel of the highest-energy electrons. In the case of phosphorus, these electrons are in a p sublevel. Therefore, phos- phorus will be found in the p-block.

Finally, use the number of valence electrons to determine the group number of the element. Phosphorus has two electrons in an s orbital and three electrons in p orbitals for a total of five valence electrons. Because there are no incomplete d or f sublevels, phosphorus must be a representative element in group 5A. A glance at the table will confirm this answer. Without using the periodic table, determine the group, period, and block in which an element with each of the following elec- tron configurations is found.

Write the electron configuration of the following elements. The larger the radius, the larger is the atom. Research shows that atoms tend to decrease in size across a period because the nuclei are increasing in positive charge while electrons are being added to sublevels that are very close in energy. As a result, the increased nuclear charge pulls the outermost electrons closer to the nucleus, making the atom smaller.

Moving down through a group, atomic radii increase. Even though the positive charge of the nucleus increases, each successive element has electrons in the next higher energy level. Electrons in these higher energy levels are located farther from the nucleus than those in lower energy levels. The increased size of higher energy level outweighs the increased nuclear charge. Therefore, the atoms increase in size. For each of the following pairs, predict which atom is larger.

Mg, Sr d. Ge, Br b. Sr, Sn e. Cr, W c. Ge, Sn Comparing elements from left to right across a period, what general trend would you predict for the energy required to remove a valence electron from an atom?

Explain the basis for your prediction. Because an electron has a negative charge, gaining electrons produces a negatively charged ion, whereas losing electrons produces a positively charged ion.

As you might expect, the loss of electrons produces a positive ion with a radius that is smaller than that of the parent atom. Conversely, when an atom gains electrons, the resulting negative ion is larger than the parent atom.

Practically all of the elements to the left of group 4A of the peri- odic table commonly form positive ions. As with neutral atoms, positive ions become smaller moving across a period and become larger moving down through a group.

These ions, although considerably larger than the pos- itive ions to the left, also decrease in size moving across a period. Like the positive ions, the negative ions increase in size moving down through a group.

Practice Problems In the following questions, the charges of ions are indicated by the superscript numbers and signs. For each of the following pairs, predict which atom or ion is larger. Explain your prediction. The first ionization energy of an element is the amount of energy required to pull the first valence electron away from an atom of the element. Atoms with high ionization energies, such as fluorine, oxygen, and chlorine, are found on the right side of the periodic table and are unlikely to form positive ions by losing electrons.

Instead, they usually gain electrons, forming negative ions. Atoms with low ionization energies, such as sodium, potassium, and strontium, lose electrons easily to form positive ions and are on the left side of the periodic table. Recall that atoms decrease in size from left to right across a period. First ionization energies generally increase across a period of elements primarily because the electrons to be removed are successively closer to the nucleus.

First ionization energies decrease moving down through a group of elements because the sizes of the atoms increase and the electrons to be removed are farther from the nucleus.

For each of the following pairs, predict which atom has the higher first ionization energy. Mg, Na d. Cl, I b. Na, Al c. Ca, Ba f. Se, Br For each of the following pairs, predict which atom forms a positive ion more easily.

Be, Ca d. K, Ca b. Sr, Sb c. Na, Si f. This principle is called the octet rule. Exceptions to this rule are hydrogen, which can gain an electron, obtaining the stable 1s2 configuration of helium, and elements in period 2, such as lithium and beryllium, that lose electrons, also obtaining the helium configuration. For example, you can predict that an element in group 6A, having a high ionization energy, will gain two electrons to achieve a stable octet configuration.

For each of the following elements, state whether it is more likely to gain or lose electrons to form a stable octet configura- tion and how many electrons will be gained or lost.

Mg Which noble-gas configuration is each of the following ele- ments most likely to attain by gaining or losing electrons? This bond involves either the transfer of electrons or sharing of electrons to varying degrees.

The nature of the bond between two atoms depends on the relative ability of each atom to attract electrons from the other, a property known as electronegativity. The maximum electronegativity value is 3. The trends in electronegativity in the periodic table are generally similar to the trends in ionization energy. The lowest electronegativity val- ues occur among the elements in the lower left of the periodic table.

These atoms, such as cesium, rubidium, and barium, are large and have few valence electrons, which they lose easily. Therefore, they have little attraction for electrons when forming a bond. These atoms are small and can gain only one or two electrons to have a stable noble-gas configuration. Therefore, when these elements form a chemical bond, their attrac- tion for electrons is large. Electronegativities generally increase across a period and decrease down through a group.

For each of the following pairs, predict which atom has the higher electronegativity. Ca, Ba b. Na, Al e. Cl, I f. Se, Br Chapter 6 Review Explain why the word periodic is applied to the table of ele- ments.

Why do elements in a group in the periodic table exhibit similar chemical properties? What chemical property is common to the elements in group 8A of the periodic table? In terms of electron configurations, what does the group number of the A-groups in the periodic table tell you? Describe the group and period trends in the following atomic properties. Describe the relationship between the electronegativity value of an element and the tendency of that element to gain or lose electrons when forming a chemical bond.

The remaining synthetic elements are created in laboratories or nuclear reactors. The elements in groups 1A through 8A of the periodic table display a wide range of properties and are called representative elements. The number of valence electrons in these elements ranges from one in group 1A to eight in group 8A, corresponding to the group numbers. The electrons are in s or p orbitals. Because elements in a given group have the same number of valence electrons, they have similar properties.

The properties are not identical, however, because the numbers of nonvalence electrons differ. For example, the ionization energy of elements in a group decreases as the atomic number increases.

For that reason, metals, which tend to lose electrons when they react, increase in reactivity as the atomic number increases. Nonmetals, which tend to gain electrons, decrease in reactivity as the atomic number increases. Such a relationship is called a diagonal relationship.

Hydrogen behaves as a metal when it loses its single electron. It behaves as a nonmetal when it gains an electron. The universe contains more than 90 percent hydrogen by mass.

Hydrogen can react explosively with oxygen. The product of this reaction is water. The main industrial use of hydrogen is in the production of ammonia. Lithium, the least reactive alkali metal, has a diagonal rela- tionship with magnesium, a group 2A metal. Sodium and potassium are the most abundant alkali metals. Potassium compounds are included in fertilizers because potassium is essential for plant growth. The remaining alkali metals—rubidium, cesium, and francium—have relatively few commercial uses.

Beryllium is used to moderate neutrons in nuclear reactors and in alloys used for non-sparking tools. Calcium is an essential element for humans, especially in maintaining healthy bones and teeth. Calcium carbonate is found in rocks such as limestone. It is used in antacids and as an abrasive. Lime, an oxide of calcium, is used to make soil less acidic, to remove pollutants from smoke- stacks, and to make mortar.

Magnesium is used in lightweight alloys and forms an oxide that is extremely heat-resistant. Sodium Na and potassium K Both sodium and potassium are in group 1A.

Sodium is above potassium in the group; thus, sodium has a higher ionization energy and is less reactive than potassium.

Compare each of the following pairs of elements in terms of group number, number of valence electrons, typical ion formed, ionization energy, and reactivity. Identify the elements. The similarity of Y to magne- sium suggests a diagonal relationship. The additional fact that Y is the least reactive element in its group make it clear that Y is lithium. Identify the following group 1A and group 2A elements based on the properties described. Many can form more than one type of ion.

The main source of boron is a compound called borax, which is used as a cleaning agent and as fireproof insulation. Aluminum, the most abundant metal, is used to make products such as cans. Gallium is used in some thermometers because it remains liquid over a wide temperature range. A branch of chemistry called organic chemistry studies the carbon compounds that control what happens in cells.

Elemental carbon occurs in various forms including the soft graphite used in pencils, and diamond, one of the hardest substances known.

Forms of an element in the same physical state that have different struc- tures and properties are called allotropes. Group 4A also contains the metalloids silicon and germanium and the metals tin and lead.

Alloys of tin such as bronze are mainly used for decorative items. Lead was one of the first metals obtained from its ore and had many uses until it was determined to be toxic.

The major current use of lead is in automo- bile storage batteries. Bacteria in soil and roots convert molecular nitrogen into compounds that can be used by plants and the animals that consume the plants. Compounds of phosphorus can be found in baking powder, cleaning products, and fertilizer. State the name of each. Then compare them in terms of group number and number of valence electrons. Identify each element as a metal, non- metal, or metalloid.

Also state a use for each element. C and Pb b. Si and P c. Ga and N a. Both C, carbon, and Pb, lead, are in group 4A and have four valence electrons. Carbon is a nonmetal. In its graphite form, it is used in pencils. Lead is a toxic metal used in automobile storage batteries. Si, silicon, is in group 4A and has four valence electrons. P, phosphorus, is in group 5A and has five valence electrons.

Phosphorus is a nonmetal. Red phosphorus provides the striking surface for matchboxes. Ga, gallium, is in group 3A and has three valence electrons. N, nitrogen, is in group 5A and has five valence electrons.

Gallium is a metal used in some thermometers; a compound of gallium is used to produce semiconductor chips used in light-powered calculators. Nitrogen is a nonmetal used to make ammonia, which is found in many cleaning products. Name the elements in each of the following pairs. Compare them in terms of group number, number of valence electrons, and metallic character.

As and Bi b. Ge and N c. B and Sn 4. An element has three valence electrons. It always loses one electron per atom when it forms an ion, and it behaves like a metal. These elements include the nonmetals oxygen and sulfur, the metalloids selenium and tellurium, and the rare radioactive metal polonium. It has two allotropes. One, O2, supports combustion and is used by organisms to release energy during respiration. The other allotrope, ozone, is an unstable, irritating gas.

Sulfur, which has ten allotropes, is used to make sulfuric acid, a compound whose production is an indication of the strength of an economy. Selenium is used in solar panels and photocopiers. Compounds of fluorine, the most electronegative element, are used in toothpaste. Chlorine is used as a disinfectant and bleach and to make certain plastics.

The silver compounds of bromine and iodine are used to coat photographic film. Iodine is important for maintaining a healthy thyroid gland. Helium is used in bal- loons and in breathing mixtures for divers. Neon and other noble gases are used to produce colored light displays. Argon, the most abundant noble gas, provides an inert atmosphere when a mixture of oxygen, heat, and sparks would be dangerous.

Argon and krypton are used to extend the life of filaments in incandescent lightbulbs. Give the name of each element. Then compare them in terms of group num- ber and number of valence electrons. State what negatively charged ion, if any, each element forms. State a use for each element.

O and Ne b. Br and Kr a. Oxygen is used by organisms to release energy during respiration. Ne, neon, is in group 8A, has eight valence electrons, and does not tend to form ions. It is used in colored light displays. A compound of bromine and silver is used to coat photographic film. Kr, krypton, is in group 8A, has eight valence electrons, and does not tend to form ions.

Krypton is used to prolong the life of filaments in incandescent lightbulbs. Practice Problems 5. Compare them in terms of group number, number of valence electrons, and typical negative ion formed if any. Also, state a use for each element. Se and Cl b. An element is a gas at room temperature. It does not form any compounds. It has eight valence electrons and is higher in atomic mass than the element phosphorus but lower in mass than arsenic. An element is metallic and radioactive.

It has six valence electrons. They are subdivided into d-block and f-block elements—the transi- tion metals and inner transition metals, respectively. The transition metals have typical metallic properties, such as malleability and electrical conductivity. Most are hard and have high melting and boiling points. Some can also lose d electrons and take on higher charges. The positive ions that have unpaired d electrons are typically colored. The movement of electrons in metals can give rise to magnetism.

If all electrons are paired, the metal is diamagnetic; that is, it is not attracted to or is slightly repelled by a magnetic field. If there is an unpaired electron, the metal is paramagnetic and is slightly attracted to a magnetic field.

The inner transition metals are placed below the main body of the periodic table. They include the lanthanide series, which are in period 6 and follow the element lanthanum, and the actinide series, which are in period 7 and follow actinium. The lanthanides have very similar properties and are silvery metals with relatively high melting points. The actinides are all radioactive, and most are synthetic elements.

They include the transuranium elements, which are elements that have an atomic number greater than Example Problem Characteristics of Transition Metals Describe the following transition metal in terms of period, block, number and pairing of d electrons, magnetic properties, hardness, and melting point: scandium Sc, atomic number Scandium tends to be paramagnetic and has only one d electron, which must be unpaired.

Compare the following pair of transition metals in terms of period number, block, number and pairing of d electrons, magnetic properties, hardness, and melting point.

Nickel Ni, atomic number 28 and yttrium Y, atomic number 39 9. For each of the following transition elements, state the period number and block, and tell whether each is a transition metal, a lanthanide, or an actinide. State what is meant by a diagonal relationship and give an example of one. Compare the alkali and alkaline earth metals in terms of posi- tion in the periodic table, number of valence electrons, and overall properties.

Compare the metallic character of the elements carbon, silicon, and lead. What do these elements have in common in terms of valence electrons and placement in the periodic table? What is meant by the term allotropes?

Give an example of allotropes. Compare the period-2 elements that are in groups 5A, 6A, 7A, and 8A in terms of number of valence electrons and reactivity. What is the difference between the transition metals and inner transition metals in terms of their final electron? Describe the placement of transition metals and inner transition metals in the periodic table and give an example of each. Chemical bonds form because of attractions between oppositely charged atoms, called ions, or between electrons and nuclei.

The outermost, or valence, electrons of atoms are the ones mainly involved in the formation of bonds. The elements within a group of the periodic table typically have the same number of valence electrons. Elements tend to react so as to achieve the stable electron con- figuration of a noble gas, typically an octet of electrons. A cation, or positive ion, is formed when an atom loses one or more electrons.

An anion, or negative ion, is formed when an atom gains one or more electrons. The periodic table is useful in predicting the charges of ions typically formed by various atoms.

Calcium Ca, atomic number 20 is an element in group 2A of the periodic table. Write the electron configuration for a neutral atom of calcium. Tell how many electrons this atom readily tends to gain or lose to form an ion. Predict the charge on the ion, write its formula, and tell whether it is a cation or an anion. Finally, write the electron configuration of this ion. A neutral atom of element 20 would have 20 electrons, giving it the electron configuration 1s22s22p63s23p64s2.

The configuration of this ion would be 1s22s22p63s23p6. Then write the formula of the ion the atom is most likely to form and identify that ion as a cation or an anion. Finally, write the electron configuration of the ion. An atom that loses one or more electrons becomes a positive ion. An atom that gains one or more electrons becomes a negative ion. The ionic bond that forms is the electrostatic force holding the oppositely charged ions together.

The total number of electrons lost must equal the total number of electrons gained. The ratio of atoms that bond ionically must therefore be such that overall electrical neutrality is maintained.

Example Problem The Formation of an Ionic Compound Atoms of magnesium Mg, atomic number 12, group 2A and chlorine Cl, atomic number 17, group 7A bond to form the ionic compound magnesium chloride. Use electron configurations and the balance of charges to determine the ratio of magnesium and chlorine atoms in magnesium chloride. A neutral atom of magnesium has the electron configuration 1s22s22p63s2, or, in abbreviated form, [Ne]3s2.

Chlorine has the electron configuration [Ne]3s23p5. Thus, there is one Mg atom for every two Cl atoms in magnesium chloride. Then determine the ratio of the atoms in the ionic compound formed in each case. The resulting structure is called a crystal lattice and contains a regular, repeating, three- dimensional arrangement of ions.

This arrangement, which involves strong attraction between oppositely charged ions, tends almost always to produce certain properties, such as high melting and boiling points and brittleness.

Ionic compounds are always non- conductors of electricity when solid but good conductors when melted. The combination of these conductivity characteristics is a very good identifier of ionic com- pounds, although each characteristic separately is not very reliable. The energy required to separate one mole of the ions of an ionic compound is called lattice energy, which is expressed as a negative quantity.

The greater that is, the more negative the lattice energy is, the stronger is the force of attraction between the ions. Lattice energy tends to be greater for more-highly-charged ions and for small ions than for ions of lower charge or large size.

For each of the following pairs of ionic compounds, state which would be expected to have the higher more negative lattice energy. LiF or KBr b. NaCl or MgS c.

MgO or RbI 8. The overall charge of any formula unit is zero.