Chem Exam recommended problems

Write the empirical formulas of the following compounds: (a) Al2Br6, (b) Na2S2O4, (c) N2O5, (d) K2Cr2O7

(a) AlBr3 (b) NaSO2 (c) N2O5 (d) K2Cr2O7

Referring to the periodic table, name (a) the halogen in the fourth period, (b) an element similar to phosphorus in chemical properties, (c) the metal in the fifth period with the lowest ionization energy, (d) an element that has an atomic number smaller than 20 and is similar to strontium.

(a) Br (b) N (c) Rb (d) Mg ( atomic number not radius)

Write the empirical formulas of the following compounds: (a) C2N2, (b) C6H6, (c) C9H20, (d) P4O10, (e) B2H6.

(a) CN (b) CH (c) C9H20 (d) P2O5 (e) BH3

Name the following compounds: (a) CdI2, (b) FeO, (c) Fe2O3, (d) TiCl4

(a) Cadmium Iodide (b) Iron(II) oxide (c) Iron(III) oxide (d) Ti(IV) chloride

Indicate which one of the two species in each of the following pairs is smaller: (a) Cl or Cl−; (b) Na or Na+; (c) O2− or S2−; (d) Mg2+ or Al3+; (e) Au or Au3+

(a) Cl is smaller because an atom gets bigger when more electrons are added (b) Na+ is smaller (c) O2- is smaller (d) Al3+ is smaller because the two ions are isoelectronic, so the ion with more protons is smaller (e) Au3+ is smaller

Identify the ions, each with a net charge of +3, that have the following electron configurations: (a) [Ar]3d3, (b) [Ar], (c) [Kr]4d6, (d) [Xe]4f14 5d6

(a) Cr3+ (b) Sc3+ (c) Rh3+ (d) Ir3+

Write the formulas for the following ionic compounds: copper bromide (containing the Cu+ ion), manganese oxide (containing the Mn3+ ion), mercury iodide (containing the (Hg2)2+ ion), magnesium phosphide

(a) CuBr (b) Mn2O3 (c) Hg2I2 (d) Mg3P2

Write the formulas for the following compounds: (a) copper(I) cyanide, (b) strontium chlorite, (c) perbromic acid, (d) hydroiodic acid, (e) disodium ammonium phosphate, (f) lead(II) carbonate, (g) tin(II) sulfite, (h) cadmium thiocyanate

(a) CuCN (b) Sr(ClO2)2 (c) HBrO4 (d) HI (e) Na2NH4PO4 (f) PbCO3 (g) SnSO3 (h) Cd(SCN)2

(a) Define the term electron affinity. (b) Explain why electron affinity measurements are made with gaseous atoms. (c) Ionization energy is always a positive quantity, whereas electron affinity may be either positive or negative. Explain

(a) Electron affinity is the energy released when an atom in the gaseous phase accepts an electron. (b) Electron affinity is a measure for atoms and molecules in the gaseous state only, since in the solid or liquid states, their energy levels would be changed by contact with other atoms or molecules. (c) The loss of one or more electrons from an atom always yields a cation, an ion with a net positive charge. Therefore ionization energy is always a positive quantity. While many first electron affinities are positive, subsequent electron affinties are always negative because considerable energy is required to overcome the repulsive forces between the electron and the negatively charged ion.

Specify the group on the periodic table in which each of the following elements is found: (a) [Ar]4s1; (b) [Ar]4s2 3d10 4p3; (c) [Ne]3s2 3p3; (d) [Ar]4s2 3d6

(a) Group 1A (b) Group 5A (c) Group 5A (d) Group 8B

Arrange the elements in each of the following groups in order of increasing electron affinity: (a) Li, Na, K; (b) F, Cl, Br, I

(a) K<Na<Li (b) I<Br<F<Cl

Give the formulas and names of the compounds formed from the following pairs of ions: (a) K+ and F−, (b) Rb+ and O2−, (c) Ba2+ and P3−, (d) Ga3+ and Se2−

(a) KF Potassium Fluoride (b) Rb2O Rubidium Oxide (c) Ba3P2 Barium Phosphide (d) Ga2Se3 Gallium Sulfide

Write the formulas for the following ionic compounds: (a) sodium oxide, (b) iron sulfide (containing the Fe2+ ion), (c) cobalt telluride (containing the Co3+ and Te2− ions), (d) barium fluoride

(a) Na2O (b) FeS (c) Co2Te3 (d) BaF2

Name the following binary molecular compounds: (a) NCl3, (b) IF7, (c) P4O6, (d) S2Cl2

(a) Nitrogen trichloride (b) Iodide heptafluoride (c) tetraphosphorus hexoxide (d) disulfur dichloride

Write chemical formulas for the following molecular compounds: (a) phosphorus tribromide, (b) dinitrogen tetrafluoride, (c) xenon tetroxide, (d) selenium trioxide

(a) PBr3 (b) N2F4 (c) XeO4 (d) SeO3

Give the formulas and names of the compounds formed from the following pairs of ions: (a) Rb+ and I−, (b) Cs+ and Se2−, (c) Sr2+ and N3−, (d) Al3+ and S2−

(a) RbI Rubidium Iodide (b) Cs2Se Cesium selenide (c) Sr3N2 strontium nitride (d) Al2S3 Aluminum sulfide

Write the formulas for the following compounds: (a) rubidium nitrite, (b) potassium sulfate, (c) sodium hydrogen sulfide, (d) magnesium phosphate, (e) calcium hydrogen phosphate, (f) potassium dihydrogen phosphate, (g) ammonium sulfate, (h) silver perchlorate

(a) RbNO2 (b) K2SO4 (c) Na2HS (d) Mg3(PO4)2 (e) CaHPO4 (f) KH2PO4 (g) (NH4)2SO4 (h) AgClO4

Determine what element is designated by each of the following: (a) fifth period, ns2 (n−1)d10 np2; (b) fourth period, ns2 (n−1)d3; (c) third period, ns2 np5; (d) sixth period, ns2.

(a) Sn (b) V (c) Cl (d) Ba

Name the following compounds: (a) NaH, (b) Li3N, (c) Na2O, (d) Na2O2

(a) Sodium hydride (b) Lithium nitride (c) Sodium oxide (d) Sodium peroxide

Write the ground-state electron configurations of the following ions, which play important roles in various biological processes: (a) Fe2+, (b) Cu2+, (c) Co2+, (d) Mn2+

(a) [Ar]3d6 (b) [Ar]3d9 (c) [Ar]3d7 (d) [Ar]3d5

Write the ground-state electron configurations of the following metal ions: (a) Ni2+, (b) Cu+, (c) Ag+, (d) Au+, (e) Au3+

(a) [Ar]3d8 (b) [Ar]3d10 (c) [Kr]4d10 (d) [Xe]4f14 5d10 e. [Xe] 4f14 5d8

Write the ground-state electron configurations of the following ions: (a) Rb+, (b) Sr2+, (c) Sn2+, (d) Te2−, (e) Ba2+, (f) In3+, (g) Tl+, (h) Tl3+

(a) [Kr], (b) [Kr], (c) [Kr]5s2 4d10, (d) [Kr]5s2 4d10 5p6, (e) [Xe], (f) [Kr]4d10 (g) [Xe]6s2 4f14 5d10 (h) [Xe] 4f14 5d10

Group the following electron configurations in pairs that would represent elements with similar properties: a. 1s2 2s2 2p5 b. 1s2 2s1 c. 1s2 2s2 2p6 d. 1s2 2s2 2p6 3s2 3p5 e. 1s2 2s2 2p6 3s2 3p6 4s1 f. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

(a) and (d), (b) and (e), (c) and (f).

Identify the following as elements or compounds: NH3, N2, S8, NO, CO, CO2, H2, SO2

(a) compound (b) element (c) element (d) compound (e) compound (f) compound (g) element (h) compound

For each of the following pairs of elements, state whether the binary compound they form is likely to be ionic or covalent. Write the empirical formula and name of the compound: (a) B and F, (b) K and Br

(a) covalent, BF3, boron trifluoride (b) ionic, KBr, potassium bromide.

For each of the following pairs of elements, state whether the binary compound they form is likely to be ionic or covalent. Write the empirical formula and name of the compound: (a) I and Cl, (b) Mg and F

(a) covalent, ICl, iodine chloride (b) ionic, MgF2, magnesium fluoride

Classify the following bonds as ionic, polar covalent, or nonpolar covalent, and explain: (a) the CC bond in H3CCH3, (b) the KI bond in KI, (c) the NB bond in H3NBCl3, (d) the CF bond in CF4

(a) covalent; the two carbon atoms are the same. (b) polar covalent; the electronegativity difference is 1.7 (c) polar covalent; the electronegativity difference is 1.0 (d) polar covalent; the electronegativity difference is 1.5

Classify the following bonds as ionic, polar covalent, or nonpolar covalent, and explain: (a) the SiSi bond in Cl3SiSiCl3, (b) the SiCl bond in Cl3SiSiCl3, (c) the CaF bond in CaF2, (d) the NH bond in NH3

(a) covalent; the two silicon atoms are the same. (b) polar covalent. The electronegativity difference is 1.2 (c) ionic. The electronegativity difference is 3.0 (d) polar covalent. The electronegativity difference is .9

How does atomic radius change (a) from left to right across a period and (b) from top to bottom in a group?

(a) decreases from left to right across a period dud to an increase in effective nuclear charge. (b) it increases from top to bottom since the orbital size increases with increasing principal quantum number.

Identify the following as elements or compounds: P4, PH3, N2O, C6H6, Hg, Cl2, N2, PCl5

(a) element (b) compound (c) compound (d) compound (e) element (f) element (g) element (H) compound

Classify each of the following as an element, a compound, or a mixture: (a) helium, (b) sugar, (c) gold, (d) hydrogen peroxide, (e) air, (f) seawater.

(a) element (b) compound (c) element (d) compound (e) mixture (f) mixture

Name the following compounds: (a) KClO, (b) Ag2CO3, (c) HNO2, (d) KMnO4, (e) CsClO3, (f) KNH4SO4, (g) Fe(BrO4)2, (h) K2HPO4

(a) potassium hypochlorite (b) silver carbonate (c) hydrogen nitrite (d) potassium permangnate (e) cesium chlorate (f) potassium ammonium sulfate (g) iron(II) perbromate (h) potassium hydrogen phosphate

Name the following compounds: (a) K3PO4, (b) CoC2O4, (c) Li2CO3, (d) K2Cr2O7, (e) NH4NO2, (f) HIO3, (g) SrSO4, (h) Al(OH)3

(a) potassium phosphate (b) cobalt oxalate (c) lithium carbonate (d) potassium dichromate (e) ammonium nitrite (f) hydrogen iodate (g) strontium sulfate (h) aluminum hydroxide

What are the functional groups?

1. carboxylic acid -COOH 2. amine -NH2 3. alchohol -OH 4. aldehyde -CHO

What is a Lewis dot symbol? What information does a Lewis dot symbol provide?

A Lewis dot symbol consists of the element's symbol surrounded by dots, where each dot represents a valence electron. The Lewis dot symbol depicts the valence electrons of atoms that interact to form compounds.

Define the term compound and explain how a compound differs from a mixture.

A compound is substance composed of two or more elements combined in a specific ratio and held together by chemical bonds. Unlike a mixture, a compound cannot be separated into simpler substances by a physical process.

Define molecular formula and empirical formula. What are the similarities and differences between the empirical formula and molecular formula of a compound?

A molecular formula shows the exact number of atoms of each element in a molecule. An empirical formula shows the lowest whole number ratio of the atoms of each element in a molecule. Similarities and differences 1. Empirical and molecular formulas are both chemical formulas. 2. The molecular formula lists all the atoms in a molecule, whereas the empirical formula shows the ratio of the number o the atoms in a molecule. 3. Empirical formulas are used to describe ionic compounds and macromolecules.

Write the outer electron configurations for the (a) alkali metals, (b) alkaline earth metals, (c) halogens, (d) noble gases

A. alkali metals: Li 2s1, Na 3s1, K 4s1, Rb 5s1, Cs 6s1, Fr 7s1 B. alkaline earth metals: Be 2s2, Mg 3s2, Ca 4s2, Sr 5s2, Ba 6s2, Ra 7s2 C. halogens: F 2s2 2p5, Cl 3s2 3p5, Br 4s2 4p5, I 5s2 5p5, At 6s2 2p5 D. Noble gases: He 1s2, Ne 2s2 2p6, Ar 3s2 3p6, Kr 4s2 4p6, Xe 5s2 5p6, Rn 6s2 6p6

Explain why alkali metals have a greater affinity for electrons than alkaline earth metals.

Alkali metals have a valence electron configuration of ns1, so they can accept another electron in the ns orbital. On the other hand, alkaline earth metals have a valence electron configuration of ns2, so they have little tendency to accept another electron unless it goes into a higher energy p orbital.

Explain the trend in electron affinity from aluminum to chlorine

Aluminum and Chlorine belong to Period 3 of the periodic table; Electron affinity increases from Al3+ to Cl-. This is because of increase in effective nuclear charge from left to right.

Define the term acid

An acid is a substance that produces hydrogen ions (H+) when dissolved in water.

What is an ionic compound? How is electrical neutrality maintained in an ionic compound?

An ionic compound consists of anions and cations. To maintain electrical neutrality, the sum of charge on the cation and the ion must be zero.

Explain why an isoelectronic series cannot include more than one member of the same group

An isoelectronic series cannot include more than one member of the same group because they would vary in electronic configuration and therefore would have different numbers of electrons.

Explain why the atomic radius of S is larger than that of O, but smaller than that of P

As we move down a group, inner electrons shield the outer electrons and quantum number n increases; the atom gets larger in moving down from O to S. As we move across a period, shielding does not change also more positive attraction is experienced by the outer electrons; the atom gets smaller in moving across from P to S.

Group the species that are isoelectronic: Be2+, F−, Fe2+, N3−, He, S2−, Co3+, Ar

Be2+ and He; Fe2+ and Co2+; N3- and F-; S2- and Ar

Describe how the naming of molecular binary compounds is different from the naming of ionic binary compounds. Explain why the two approaches are different

Binary molecular compounds are named according to the number of atoms in the compound. Each number has its own prefix such as mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, etc. Ionic compounds do not have prefixes but the anion always has the suffix -ide The reason why the prefix is not needed to specify the number of ions in an ionic compound is because the sum of the charges of the two atoms is zero. This is why the two approaches are different.

Specify which of the following elements you would expect to have the greatest electron affinity: He, K, Co, S, Cl

Cl would be expected to have the highest electron affinity, because the tendency to accept electrons increases as we move from right to left across a period.

Arrange the following atoms in order of increasing atomic radius: Si, Mg, Cl, P, Al

Cl<P<Si<Al<Mg

Explain which of the following cations is larger and why: Cu+ or Cu2+

Cu+ is larger because it has one more electron.

The first and second ionization energies of K are 419 and 3052 kJ/mol, and those of Ca are 590 and 1145 kJ/mol, respectively. Compare their values and comment on the differences.

Effective nuclear charge increases across a row, so we would expect the first ionization energy for calcium to be higher. For potassium, however, the second electron must come from the atom's noble gas core, which accounts for the much higher second ionization energy.

Explain the term effective nuclear charge

Effective nuclear charge is the actual magnitude of positive charge that is experienced by an electron in the atom. In atoms with more than one electron, the electrons are simultaneously attracted to the nucleus and repelled by one another.

Define electronegativity and explain the difference between electronegativity and electron affinity. Describe in general how the electronegativities of the elements change according to their position in the periodic table

Electronegativity is the ability of an atom in a compound to draw electrons to itself that it shares in a chemical bond with another atom, whereas electron affinity refers to an isolated atom's ability to attract an additional electron in the gas phase. Electron affinity is an experimentally measurable quantity, whereas electronegativity is an estimated number that cannot be measured directly. In general, electronegativity increases from left to right across a period in the periodic table. Within a group, electronegativity decreases with increasing atomic number.

Describe the general layout of a modern Periodic table

Elements are arranged in the periodic table based on the type of subshell containing the outermost electrons-the main group elements, the noble gases and the transition elements.

What is the most important relationship among elements in the same group in the periodic table?

Elements in the same group have similar physical and chemical properties.

A metal ion with a net +3 charge has five electrons in the 3d subshell. Identify the metal.

Fe3+

Explain how the octet rule applies to hydrogen

For hydrogen, the octet rule dictates that it has two electrons, giving it the electronic configuration of the nearest noble gas, helium.

Explain the concept of formal charge

Formal charge is a way of keeping track of the valence electrons in a species. Formal charge on an atom in a Lewis structure is equal to the number of valence electrons (group number) minus half of the electrons it shares with other atoms in the structure. Formal charge is determined by comparing the number of electrons associated with an atom in a Lewis structure with the number of electrons that would be associated with the isolated atom.

Compare single, double, and triple bonds in a molecule, and give an example of each. For the same bonding atoms, how does the bond length change from single bond to double bond to triple bond?

If a bond is formed between two atoms and consists of one shared pair of electrons, it is called a single bond. A double bond is a multiple bond in which the atoms share two pairs of electrons; in a triple bond, atoms share three pairs of electrons. Single bond example is ethane (C2H6) C-C Double bond example is ethylene (C2H4) C=C Triple bond example is acetylene (C2H2) C≡C Single bonds are longer than double bonds, and double bonds are longer than triple bonds.

Define the terms cation and anion.

If an atom loses one or more electrons, it has a net positive charge and is known as a cation. If an atom gains electrons, it has a net negative charge and is known as anion.

Explain why, for a cation and an anion that are isoelectronic, the anion is larger than the cation.

In an isoelectronic series, the anions are larger than the cations because they have smaller nuclear charges.

Describe the difference between naming HF when it is in the gas phase and HF when it is dissolved in water. Give another example of a compound that has one name when in the gas phase and another name when dissolved in water

In the gas phase, HF is hydrogen fluoride; when dissolved in water, it is called hydrofluoric acid. Another example of a compound that has one name in the gas phase and another name when dissolved in water is HCl. In the gaseous phase this is known as hydrogen chloride.; when dissolved in water it is known as hydrochloric acid.

Explain what ionic bonding is.

Ionic bonding refers to the electrostatic attraction between oppositely charged ions held together in an ionic compound.

Define ionic radius. How does the size of an atom change when it becomes (a) an anion and (b) a cation?

Ionic radius is the radius of cation or an anion (a) when an atom gains one or more electrons and becomes an anion, its radius increases due to increased electron-electron repulsions (b) when an atom loses an electron and becomes a cation, its radius decreases due in part to a reduction in electron-electron repulsions in the valence shell.

Define ionization energy. Explain why ionization energy measurements are usually made when atoms are in the gaseous state. Why is the second ionization energy always greater than the first ionization energy for any element? What types of elements have the highest ionization energies and what types have the lowest ionization energies?

Ionization energy is the minimum energy required to remove an electron from an atom in the gas phase. Ionization energy measurements are usually made when atoms are in the gaseous state, since in the solid or liquid state, their energy levels would be changed by contact with other atoms or molecules. The second IE for any element will always be greater than the first IE because of ever-increasing amounts of energy-that is, it is harder to remove an electron from a cation than from a anion. As effective nuclear charge increases, IE also increases from left to right of the periodic table. Usually non-metals have high ionization energy and metals have the lowest ionization energy.

Explain why ions with charges greater than ±3 are seldom found in ionic compounds

Ions with charges greater than ±3 are seldom found in ionic compounds because the fourth ionization energy too high and the fourth electron affinity energy is too low to allow ionic compounds to be formed.

Arrange the following in order of increasing first ionization energy: F, K, P, Ca, and Ne.

K<Ca<P<F<Ne

What is lattice energy and what does it indicate about the stability of an ionic compound?

Lattice energy is the amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase. The greater the lattice energy, the more stable the compound.

What is a main group element? Give names and symbols of six main group elements: two metals, two nonmetals and two metalloids

Main group elements are those in Groups 1A through 7A. Examples of symbols of six main group elements are: Metals Sodium (Na) Calcium (Ca) Nonmetals Nitrogen (N) Sulfur (S) Metalloids Germanium (Ge) Antimony (Sb)

List the properties of metals.

Metals tend to be shiny, lustrous, malleable, ductile, and conducting (for both heat and electricity). They typically lose electrons to form cations, and they tend to form ionic compounds.

Classify each of the following as a metal, a nonmetal, or a metalloid: Sb, Kr, Co, Na, Al, F, Sr, As, Br, Ge.

Metals: Co, Na, Sr, Al Nonmetals: F, Br, Kr Metalloids: As, Ge, Sb

List the following ions in order of increasing ionic radius: N3−, Na+, F−, Mg2+, O2−

Mg2+ < Na+ < F- < O2- < N3- because in an isoelectronic series the ion with the fewest protons will have the largest ionic radius, and the ion with the most protons will have the smallest ionic radius.

Explain the difference between the terms molecular mass and formula mass. To what type of compound does each term refer?

Molecular mass is used for molecules that are discrete (covalent compounds). Formula mass refers to the mass of the empirical formula of an ionic compound.

Which is the largest atom in the third period of the periodic table

Na (Sodium) is the largest atom in the third period.

Do formal charges represent an actual separation of charges?

No

List the properties of nonmetals.

Nonmetals tend to be brittle and not good conductor (for either hear or electricity). They can gain electrons to form anions but they commonly form molecular compounds.

Which is the smallest atom in Group 6A

O (Oxygen) is the smallest atom in Group 6A

Arrange in increasing order of size for the following atoms: Na, Mg, O, and S

O<S<Mg<Na

What is the difference between an inorganic compound and an organic compound?

Organic compounds contain carbon and hydrogen, sometimes in combination with other elements such as oxygen, nitrogen, sulfur, and the halogens. Inorganic compounds generally do not contain carbon, although some carbon-containing species are considered inorganic.

Arrange in increasing order of size for the following atoms: K, Ca, S, and Se

S<Se<Ca<K

Use the first-row transition metals (Sc to Cu) as an example to illustrate the characteristics of the electron configurations of transition metals.

Sc 3d14s2, Ti 3d2 4s2, V 3d3 4s2, Cr 3d5 4s1, Mn 3d5 4s2, Fe 3d6 4s2, Co 3d7 4s2, Ni 3d8 4s2, Cu 3d10 4s1 Transitional metals either have incompletely filled d subshell or readily lose electrons to achieve incompletely filled d subshells. For the ions of the transition elements the 3d orbital are lower in energy than the 4s orbitals. Therefore, the electrons most easily lost are those in the outermost principal level, the ns.

An anion with a net −2 charge has a total of 36 electrons. Identify the element from which the anion is derived

Se; Selenium since it originally has 34 electrons, but the -2 charge means that it gains 2 electrons.

An atom of a certain element has 16 electrons. Consulting only the periodic table, identify the element and write its ground-state electron configuration.

Sulfur, S. The electron configuration is 1s2 2s2 2p6 3s2 3p4 or [Ne]3s2 3p4.

Define atomic radius and explain why such a definition is necessary

The atomic radius is the distance between the nucleus of an atom and its valence shell. Metallic radius is half the distance between nuclei of two adjacent, identical metal atoms. Covalent radius is half the distance between adjacent identical nuclear that are connected by a chemical bond. These definitions are necessary to avoid the intuitive of the atomic radius. According the the quantum mechanical model of the atom, there is no specific distance from the nucleus beyond which an electron may not be found. Therefore, atomic radius requires a specific definition

What does a chemical formula represent? Determine the ratio of the atoms in the following molecular formulas: (a) NO, (b) NCl3, (c) N2O4, (d) P4O6

The chemical formula denotes the composition of the substance. (a) 1:1 (b) 1:3 (c) 2:4 = 1:2 (d) 4:6 = 2:3

Explain why the chemical formulas of ionic compounds are usually the same as their empirical formulas

The chemical formulas of ionic compounds are generally empirical formulas because an ionic compound consists of a vast array of interspersed cations and anions called a lattice, not discrete molecular units.

Why is the radius of the lithium atom considerably larger than the radius of the hydrogen atom?

The electron configuration of lithium is 1s2 2s1. The two 1s electrons shield the 2s electron effectively from the nucleus. Consequently, the lithium atom is considerably larger than the hydrogen atom.

In general, the first ionization energy increases from left to right across a given period. Aluminum, however, has a lower first ionization energy than magnesium. Explain.

The group 3A elements (such as Al) all have a single electron in the outermost p subshell, which is well shielded from the nuclear charge by the inner electrons and the ns2 electrons. Therefore, less energy is needed to remove a single p electron than to remove a pair s electron from the same principal energy level (such as for Mg)

State the laws of definite proportions and multiple proportions. Illustrate each with an example

The law of definite proportions states that different samples of a given compound always contain the same elements in the same mass ratio. The law of multiple proportions states that if two elements can combine to form two or more different compounds, the masses of one element that combine with a fixed mass of the other element can be expressed in ratios of small whole numbers

Two atoms have the electron configurations 1s22s22p6 (Ne) and 1s22s22p63s1 (Na). The first ionization energy of one is 2080 kJ/mol, and that of the other is 496 kJ/mol. Match each ionization energy with one of the given electron configurations. Justify your choice.

The lone electron in the 3s orbital will be shielded from the nuclear charge by the filled inner shell, so the IE of 496 kJ/mol would would be paired with 1s22s22p63s1 (Na). The 1s22s22p6 (Ne) is a noble gas electron configuration. Therefore, it will be extremely difficult to remove an electron. the 2p electron is also no shielded by electron in the same energy level. 1s22s22p6(Ne) is paired with the IE 2080 kJ/mol

How does the electron configuration of ions derived from main group elements give them stability?

The main group elements tend to form ions that have half-filled or completely filled valence shells.

What is the molar mass of a compound? What are the commonly used units for molar mass?

The molar mass of a substance is the mass of one mole of a substance. Usually molar mass is expressed in grams per mole.

What is meant by the term molecular mass, and why is the molecular mass that we calculate generally an average molecular mass?

The molecular mass is the mass in atomic mass units of an individual molecule. It uses the molecular formula and the atomic masses from the periodic table, which are average atomic masses.

Explain how ionization energy and electron affinity determine whether atoms of elements will combine to form ionic compounds

The net energy charge associated with the formation of the ions must be less than the energy released when the electrostatic attraction resulting from the cation and anion draws them together to form an ionic compound.

What are valence electrons?

The outermost electrons of an atom. They determine how atoms interact with each other.

Explain the octet rule and why it applies mainly to second period elements

The valence subshells of first-period elements can hold only two electrons, but in second-period elements the subshells can hold a total of eight electrons. So, it is mainly second-period elements that lose, gain, or share electrons attain the noble gas configuration. The octet rule can be violated by higher period elements expanding their valence shells into empty d orbitals.

What do we mean when we say that two ions or an atom and an ion are isoelectronic?

Their electron configurations are the same.

Rank the following ions in order of increasing number of unpaired electrons: Ti3+, Fe2+, V3+, Cu+, Mn4+ hint (fill one electron in each orbital)

Ti3+ 1 unpaired electron Fe2+ 4 unpaired electrons V3+ 2 unpaired electrons Cu+ 0 unpaired electrons Mn4+ 3 unpaired electrons

Describe how the knowledge of the percent composition by mass of an unknown compound can help us identify the compound

Using the percent composition by mass, the empirical formula can be determined. Find the empirical formula mass by adding the molar masses of the elements. Divided the molar mass by the empirical formula mass to get a whole number. Multiply the result by each subscript in the empirical formula to determine the molecular formula of the compound.

For a given pair of bonded atoms, explain how bond length relates to bond strength

When bond length decreases, the bond strength increases.

Explain why there is a greater increase in effective nuclear charge from left to right across a period than there is from top to bottom in a group

While moving from left to right across a period, the effective nuclear charge increases steadily because the number of core electrons remains the same. But when moving from top to bottom in a group, the valence electrons are effectively shielded by the core electrons. So there is a greater increase in effective nuclear charge from left to right across a period than there is from top to bottom in a group.

What is a polar covalent bond? Name two compounds that contain one or more polar covalent bonds

a bond in which electrons are shared but are not shared equally by the two atoms. H2O and HCl are two examples of compounds that contain polar covalent bonds.

On the basis of their positions in the periodic table, select the atom with the larger atomic radius in each of the following pairs: a. Mg, P; b. Sr, Be; c. As, Br; d. Cl, I; e. Xe, Kr.

a. Mg is larger. It is to the left of P in Period 3 b. Sr is larger. It is below Be is Group 2A c. As is larger. It is to the left of Br in Period 4 d. I is larger. It is below Cl in Group 7A e. Xe is larger. It is below Kr in Group 8A