#31 Key
•Download as DOC, PDF•
3 likes•19,674 views
1) Metal character increases down a group and decreases across a period on the periodic table. Metals have low ionization energies and form cations by losing electrons. 2) Most metal oxides are basic, while most nonmetal oxides are acidic. When nonmetal oxides react with water, acids are formed. 3) The octet rule states that atoms bond to achieve a full outer shell of 8 electrons. Atoms gain, lose or share electrons to achieve an octet configuration.
Report
Share
Report
Share
![KEY
GENERAL CHEMISTRY-I (1411)
S.I. # 31
1. Metal characteristic increases _______B_______ a group.
a. up b. down c. across
2. Metal character decreases _______C_______ a period.
a. up b. down c. across
3. Metals have ______B________ ionization energies.
a. high b. low c. zero d. positive e. negative
4. Group 1A metals for _____M+____ Ions and group 2A form ____M2+_____ions.
5. Most metal oxides are _____B______
a. acidic b. basic c. neutral
6. Most non metal oxides are: ____A______
a. acidic b. basic c. neutral
7. When nonmetal oxides react with water what is formed? An Acid
8. Complete and balance the following equations.
P4O10(s) + H2O(l) 4H3PO4(aq)
CO2(g) + H2O(l) H2CO3(aq)
9. What are some properties of metalloids?
Metalloids have properties that are intermediate between metals and nonmetals.
Metalloids have found fame in the semiconductor industry.
10. Using only the periodic table, arrange the following atoms in increasing radius
(a) Cs, K, Rb K < Rb < Cs__________________
(b) In, Te, Sn Te < Sn < In_________________
(c) P, Cl, Sr Cl < P < Sr___________________
11. What is the octet rule?
Atoms will gain, lose or share electrons to achieve the nearest noble gas electron configuration.
Except for H and He, this corresponds to eight electrons in the valence shell, thus the term octet rule
12. How many electrons must a sulfur atom gain to achieve an octet in its valence
shell? S: [Ne]3s23p4 A sulfur atom has six valence electrons, so it must gain two
electrons to achieve an octet.
13. If an atom has the electron configuration 1s22s22p3, how many electrons must it
gain to achieve an octet?
1s22s22p3 = [He]2s22p3 The atom (N) has five valence electrons and must gain three
electrons to achieve an octet](https://image.slidesharecdn.com/chem-i-si-31key-1272476054-phpapp01/85/31-Key-1-320.jpg)
![KEY
14. a. Write the electron configuration for the element scandium, Sc.
Sc: 1s22s22p63s23p64s23d1 = [Ar]4s23d1
b. How many valence electrons does this atom possess?
Scandium has three (3) valence electrons.
c. What distinguishes these valence electrons from the other electrons in the
atom?
These valence electrons are available for chemical bonding, while the core
electrons do not participate in bonding.
15. Use Lewis symbols to represent the reaction that occurs between Mg and Br
atoms.
. . . . . .
. . . .
Mg +. Br . +. Br . . Mg2+ + 2[ .Br. ] -
. . . . . . .
16. What ionic compound is expected to form from combination of the following
pairs of elements?
a. barium and oxygen BaO___________________
b. rubidium and iodine RbI ___________________
c. lithium and sulfur Li2S ___________________
d. bromine and magnesiumMgBr2___________________
17. Which of these elements is unlikely to form covalent bonds: S, H, K, Ar, Si?
Explain your choices
K and Ar. K is an active metal with one valence electron . It is most likely to achieve an octet by losing
this single electron and to participate in ionic bonding. Ar has a stable octet of valence electrons; it is not
likely to form chemical bonds of any type.
18. a.
What is the most electronegative element in group 6A? O_____
b.
Which is most electronegative out of Al, Si, P? P_____
c.
Which is most electronegative out of Ga, P, Cl, Na? Cl_____
d.
The element most likely to form an ionic compound with Ba out of
K, C, Zn, F F
19. Arrange the bonds in each of the following sets in order of increasing polarity:
a. C--F, O—F, Be—F O—F < C—F < Be—F
b. O—Cl, S—Br, C—P S—Br < C—P < O—Cl
c. C—S, B—F , N—O C—S < N—O < B—F
20. Does the octet rule apply to ionic as well as covalent compounds? Explain.
The octet rule applies to the individual ions in an ionic compound. A cation loses electrons and the anion
gains electrons to achieve an octet. For example: in MgCl2, Mg loses 2 e- to become Mg2+ with the
electron configuration of [Ne]. Each Cl atom gains one electron to form Cl- with the electron configuration
of [Ar].
21. Draw the Lewis structures for each of the following molecules or ions. Which
do not obey the octet rule?](data:image/gif;base64,R0lGODlhAQABAIAAAAAAAP///yH5BAEAAAAALAAAAAABAAEAAAIBRAA7)
Recommended
SI #5 Key
This document provides a summary of key concepts in general chemistry:
1. It defines atomic number as the number of protons in an atom and mass number as the number of protons plus neutrons. The mass number can vary without changing elemental identity.
2. It provides the number of protons, neutrons, and electrons for various atoms and isotopes.
3. It identifies various elements as metals, metalloids, or nonmetals and provides their names.
4. It predicts the charges of ions for various elements.
5. It predicts the chemical formulas and names of compounds formed by combinations of elements.
#31
This document contains a chemistry practice test with multiple choice and free response questions about periodic trends, ion formation, Lewis structures, the octet rule, and electronegativity. Some key points covered are:
1. Metal character increases down a group and decreases across a period on the periodic table. Metals have low ionization energies.
2. Group 1A metals form 1+ ions and group 2A metals form 2+ ions. Most metal oxides are basic while most nonmetal oxides are acidic.
3. The octet rule states that main group elements gain, lose, or share electrons to achieve an octet of 8 electrons in their valence shell. Not all molecules obey this rule
#30 Key
The document provides information about trends in the periodic table including:
1) Ion size increases down a column due to increasing nuclear charge.
2) Elements with the same number of electrons form isoelectronic series.
3) Ionization energy decreases down a group and increases across a period as nuclear charge increases.
4) Reactivity of alkali metals increases down a group as electrons are more readily removed further from the nucleus.
#33
This document contains two chemistry questions. The first question asks about the similarities and differences between atomic orbitals and molecular orbitals, why the bonding molecular orbital of H2 is at a lower energy than in a hydrogen atom, and how many electrons can be placed in each molecular orbital. The second question asks about the relationships between bond order, bond length, and bond energy, and whether Be2 or Be2+ would be expected to exist according to molecular orbital theory.
#32 Key
The document summarizes key concepts in general chemistry including:
1) Tetrahedral molecules like methane and perchlorate ion have bond angles of 109.5 degrees. Trigonal pyramidal ammonia requires specifying two parameters - the H-N-H bond angle and N-H bond distance.
2) The number of electron domains is determined by counting bonds and lone pairs. Bonding domains contain bonding electrons while nonbonding domains contain lone pairs.
3) Molecules with 3, 4, 5, or 6 electron domains exhibit trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral electron domain geometries respectively.
4) Electron domain angles of
Ions and bonding
teachers answers for teaching ions and bonding,
spaces provided for students to copy down diagrams and drawings of ions' electronic configuration and covalent bonding between molecules
C3 syllabus statements
This document outlines the syllabus statements for a course on atoms, elements, compounds and chemical bonding. It covers physical and chemical changes, the differences between elements, compounds and mixtures, atomic structure and the periodic table, ions and ionic bonds, molecules and covalent bonds, and giant covalent structures. Key concepts include the structure of the atom, electron configuration, isotopes, ion formation, ionic and covalent bonding, and drawing dot-and-cross diagrams to represent molecules.
Ions and bonding student copy
This document provides information about ions and bonding. It discusses how atoms gain or lose electrons to achieve stable octet configurations, forming ions. It explains that metals donate electrons to form positive ions, while non-metals gain electrons to form negative ions. Ionic bonds form when oppositely charged ions attract. The document also covers covalent bonding between non-metal atoms by shared electron pairs. Students are asked to determine ion charges, draw ionic compounds and molecules, and identify concepts like electronic configurations and bonding types.
Recommended
SI #5 Key
This document provides a summary of key concepts in general chemistry:
1. It defines atomic number as the number of protons in an atom and mass number as the number of protons plus neutrons. The mass number can vary without changing elemental identity.
2. It provides the number of protons, neutrons, and electrons for various atoms and isotopes.
3. It identifies various elements as metals, metalloids, or nonmetals and provides their names.
4. It predicts the charges of ions for various elements.
5. It predicts the chemical formulas and names of compounds formed by combinations of elements.
#31
This document contains a chemistry practice test with multiple choice and free response questions about periodic trends, ion formation, Lewis structures, the octet rule, and electronegativity. Some key points covered are:
1. Metal character increases down a group and decreases across a period on the periodic table. Metals have low ionization energies.
2. Group 1A metals form 1+ ions and group 2A metals form 2+ ions. Most metal oxides are basic while most nonmetal oxides are acidic.
3. The octet rule states that main group elements gain, lose, or share electrons to achieve an octet of 8 electrons in their valence shell. Not all molecules obey this rule
#30 Key
The document provides information about trends in the periodic table including:
1) Ion size increases down a column due to increasing nuclear charge.
2) Elements with the same number of electrons form isoelectronic series.
3) Ionization energy decreases down a group and increases across a period as nuclear charge increases.
4) Reactivity of alkali metals increases down a group as electrons are more readily removed further from the nucleus.
#33
This document contains two chemistry questions. The first question asks about the similarities and differences between atomic orbitals and molecular orbitals, why the bonding molecular orbital of H2 is at a lower energy than in a hydrogen atom, and how many electrons can be placed in each molecular orbital. The second question asks about the relationships between bond order, bond length, and bond energy, and whether Be2 or Be2+ would be expected to exist according to molecular orbital theory.
#32 Key
The document summarizes key concepts in general chemistry including:
1) Tetrahedral molecules like methane and perchlorate ion have bond angles of 109.5 degrees. Trigonal pyramidal ammonia requires specifying two parameters - the H-N-H bond angle and N-H bond distance.
2) The number of electron domains is determined by counting bonds and lone pairs. Bonding domains contain bonding electrons while nonbonding domains contain lone pairs.
3) Molecules with 3, 4, 5, or 6 electron domains exhibit trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral electron domain geometries respectively.
4) Electron domain angles of
Ions and bonding
teachers answers for teaching ions and bonding,
spaces provided for students to copy down diagrams and drawings of ions' electronic configuration and covalent bonding between molecules
C3 syllabus statements
This document outlines the syllabus statements for a course on atoms, elements, compounds and chemical bonding. It covers physical and chemical changes, the differences between elements, compounds and mixtures, atomic structure and the periodic table, ions and ionic bonds, molecules and covalent bonds, and giant covalent structures. Key concepts include the structure of the atom, electron configuration, isotopes, ion formation, ionic and covalent bonding, and drawing dot-and-cross diagrams to represent molecules.
Ions and bonding student copy
This document provides information about ions and bonding. It discusses how atoms gain or lose electrons to achieve stable octet configurations, forming ions. It explains that metals donate electrons to form positive ions, while non-metals gain electrons to form negative ions. Ionic bonds form when oppositely charged ions attract. The document also covers covalent bonding between non-metal atoms by shared electron pairs. Students are asked to determine ion charges, draw ionic compounds and molecules, and identify concepts like electronic configurations and bonding types.
Ejercicio de enlaces y perioricidad
1. The Group 8A elements are inert gases because they have stable electron configurations and do not participate in chemical reactions.
2. Atoms are stable when they resist changes in their electron arrangements.
3. When sodium reacts with chlorine, each sodium atom loses 1 electron.
Transition metals
This document discusses the nomenclature and isomers of coordination complexes. It begins by defining ligands and describing systems for naming complexes based on the central metal ion, ligands, and ligand coordination. It then describes three types of isomers that can occur in complexes: structural isomers which have different ligands, linkage isomers which have the same ligands bonded differently, and stereoisomers which have the same ligands in different spatial arrangements. The document goes on to describe crystal field theory and ligand field theory, which explain the splitting of d-orbital energies in an octahedral ligand field based on ligand type. Strong-field ligands create a large energy gap between t2g and eg* orbitals, while weak-field ligands create
Adv chem chapt 9
The molecular orbital model can be used to predict molecular structure, bonding properties, and magnetism. For diatomic molecules, atomic orbitals combine to form molecular orbitals through constructive or destructive interference. Bonding molecular orbitals are lower in energy than antibonding orbitals. Electron configuration and bond order indicate bond strength and stability. Heteronuclear diatomics form molecular orbitals between different atomic orbitals, while delocalized π bonding is seen in resonant structures like benzene.
Metal carbonyls
Metal carbonyls are coordination complexes of transition metals with carbon monoxide ligands. They were first synthesized in 1868 by passing carbon monoxide over platinum. Metal carbonyls typically obey the 18 electron rule and are often diamagnetic. They have applications as catalysts in organic synthesis and in producing pure metals like nickel. Precautions must be taken when using metal carbonyls due to their toxicity.
Naming ionic compounds writing emp formulas
This document discusses how to write empirical formulas for ionic compounds. It explains that ionic compounds are electrically neutral, meaning the positive charge of cations must cancel out the negative charge of anions. It provides an example of determining the charges on potassium (K+) and bromine (Br-) ions and writing the formula for the ionic compound potassium bromide (KBr). The document also gives a mini quiz example of determining the charges on lithium (Li+) and fluorine (F-) ions and writing the formula for the ionic compound lithium fluoride (LiF).
Chemical Bonding and Formula Writing
The document discusses chemical bonding and formula writing. It explains that there are two main types of bonds: ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons between atoms to form ions, while covalent bonds involve the sharing of electrons between atoms. It provides examples of common ionic compounds and their formulas, such as NaCl, as well as explaining how to write formulas for compounds containing polyatomic ions like OH-. Covalent bonding is demonstrated through the example of methane, CH4, where carbon shares electrons with four hydrogen atoms to fill their outer energy levels.
Naming ionic compounds intro
This document discusses the formation of ionic bonds between ions. Ionic bonds are electrostatic attractions between oppositely charged ions. Cations are positively charged ions formed when metals lose electrons. Anions are negatively charged ions formed when nonmetals gain electrons. Ionic bonds generally form between metallic cations and nonmetallic anions. The document provides examples of common cations and anions and identifies metals, nonmetals, and metalloids on the periodic table.
Chemistry - Chp 9 - Chemical Names and Formulas - Notes
1) This document provides information on naming and writing formulas for ionic and molecular compounds. It discusses monatomic and polyatomic ions, and how to determine the charges on cations and anions.
2) Rules are provided for writing formulas and naming ionic compounds, including those containing transition metals. Prefixes are introduced for naming molecular compounds based on the number of atoms present.
3) Examples are worked through for writing formulas from names and names from formulas for different types of compounds. Practice problems are also included for students to try.
2012 ppt unit 2 3 ionic bonding djy r1
Ionic bonding results from the electrostatic attraction between oppositely charged ions. When a metal atom reacts with a nonmetal, the metal typically loses electrons to form a cation while the nonmetal gains electrons to form an anion. These oppositely charged ions are then attracted to each other, forming an ionic bond. Ionic compounds have crystalline structures where the ions are arranged in repeating patterns. Their strong ionic bonds make them brittle with high melting points. Many ionic compounds dissolve in water to form electrolyte solutions where the ions are free to move.
Chemistry - Chp 9 - Chemical Names and Formulas - PowerPoint
This document covers naming and writing formulas for ions, ionic compounds, molecular compounds, acids, and bases. It provides objectives and definitions for each section, examples of naming and writing formulas, and explanations of naming conventions and rules including prefixes, charges, and endings for different types of compounds.
Naming & Writing Ionic Formulas PowerPoint
The document provides information on naming ionic compounds and writing their formulas. It discusses the rules for writing formulas, such as ensuring the total positive and negative charges are equal. It describes the "swoosh and reduce" method for determining formulas, which involves writing the symbols and charges of the elements and crossing the charges to become subscripts on the other element. The document also covers naming cations and anions and how to combine them to name complete ionic compounds.
Back donation
This presentation discusses back-donation of CO ligands to transition metals. Back-donation occurs when electrons from the metal's d-orbitals are donated into the CO ligand's anti-bonding π* orbital, weakening the carbon-oxygen bond. CO can bind to single metals through both sigma and pi bonding. It can also bridge between multiple metal atoms, further weakening the carbon-oxygen bond. Infrared spectroscopy shows that as more metals are involved in back-donation to a CO ligand, its carbon-oxygen stretching frequency decreases due to the bond being more weakened.
Theory of coordination_compounds[1]
The document discusses the theory of coordination compounds. It defines coordination compounds as products of Lewis acid-base reactions where ligands bond to a central metal atom via coordinate covalent bonds. It describes coordination bonds and terminology like ligands, central metal ion, oxidation state, and coordination number. The document also discusses Werner's theory of coordination complexes, valence bond theory, types of ligands, nomenclature of coordination compounds, and limitations of Werner's theory.
transition element
Transition metals have several characteristic properties including high melting points, hardness, and electrical conductivity. They are found in the d-block of the periodic table. The colors of transition metal compounds are determined by the energies absorbed during electron transitions between d orbitals. Spin allowed transitions that don't change the number of paired/unpaired electrons absorb lower energies, appearing as red/yellow colors. Spin forbidden transitions changing this pairing absorb higher energies, appearing as blue/green colors. Coordination number and bond strength also impact the energies absorbed and thus the colors.
Valence Bond Theory
Valence Bond Theory (VBT) explains the nature of metal-ligand bonding in complex compounds. VBT assumes that the central metal atom forms hybrid orbitals with its outer shell orbitals that overlap with ligand orbitals to form sigma bonds. There are two types of hybridization that result in octahedral geometry: d2sp3 for "inner-orbital" complexes with paired electrons, and sp3d2 for "outer-orbital" complexes with unpaired electrons. Examples that form via d2sp3 hybridization, making them inner-orbital complexes, include ferricyanide ([Fe(CN)6]3-), ferrocyanide ([Fe(CN)6]4-
Chapter 6 Notes
Chemical bonds form between atoms to achieve more stable arrangements with lower potential energy. The main types of bonds are ionic, covalent, and metallic. Ionic bonds form between oppositely charged ions, covalent bonds form when atoms share electrons, and metallic bonds result from delocalized electrons being shared among many atoms in a metal. The degree of ionic or covalent character in a bond depends on the electronegativity difference between the atoms. Molecular compounds are made of molecules held together by covalent bonds, while ionic compounds are made of ionic bonds between cations and anions.
Formula writing jhe
This document discusses the steps and rules for writing chemical formulas. It begins by outlining the objectives of learning to write formulas and defining key terms like subscripts and coefficients. It then explains the four main steps to write formulas: 1) Determine element symbols 2) Predict oxidation states 3) Write symbols with charges 4) Balance charges. The rest of the document outlines four main rules for writing subscripts and parentheses correctly based on oxidation states and whether elements appear only once or multiple times in a compound.
Bonding in coordination compound(werners theory)
This document discusses Werner's theory of coordination compounds and bonding in coordination compounds. According to Werner's theory, metal atoms in coordination compounds have both primary and secondary valencies. Primary valencies are ionizable and satisfy the compound's oxidation state, while secondary valencies are non-ionizable and satisfy the compound's coordination number through coordinate covalent bonds with electron pair donors like ligands. The document also discusses Sidgwick's effective atomic number rule and how the valence bond theory explains the geometry, hybridization, and magnetic properties of coordination compounds.
Lecture 9.1 & 9.2- Naming Salts
The document discusses chemical names and formulas of ions. It explains that monatomic cations form when metals lose electrons and gain a positive charge equal to their group number. Monatomic anions form when non-metals gain electrons and have a negative charge equal to the group number minus 8. Polyatomic ions consist of multiple atoms and have an overall positive or negative charge. Transition metals can form multiple ions with different charges that must be specified. Ionic compounds are named by writing the cation first followed by the anion.
Ch 8 Ionic Compounds
1. Ionic compounds form when metals react with nonmetals, resulting in the transfer of electrons and formation of cations and anions bonded ionically.
2. Ionic compounds have high melting points and boiling points due to the strong electrostatic forces between cations and anions in the crystalline lattice structure.
3. Metallic bonds form when metal cations attract delocalized electrons in the "sea of electrons" that moves throughout the entire metallic crystal. This produces metallic properties like conductivity and ductility.
#12 Key
1. The document defines key chemistry terms including: aqueous solutions, solvents, solutes, electrolytes, nonelectrolytes, strong/weak electrolytes, and solvation.
2. Precipitation reactions are defined as reactions where an insoluble product called a precipitate forms. Molecular, complete ionic, spectator ions, and net ionic equations are also defined.
3. Strong acids and bases are listed as well as examples of soluble and insoluble compounds in water. Spectator ions are identified for sample precipitation reactions.
#22 Key
1. An object can possess energy through kinetic energy from its motion or potential energy from its position. Kinetic energy depends on mass and velocity while potential energy depends on position.
2. A kilowatt-hour equals 3.6 million joules. An adult radiates heat at the same rate as a 100-watt bulb, so in 24 hours they radiate about 2.1 kilocalories of heat energy.
3. Chemical reactions that release heat are exothermic, while those that absorb heat are endothermic. Heat change equals internal energy change during constant pressure processes.
More Related Content
What's hot
Ejercicio de enlaces y perioricidad
1. The Group 8A elements are inert gases because they have stable electron configurations and do not participate in chemical reactions.
2. Atoms are stable when they resist changes in their electron arrangements.
3. When sodium reacts with chlorine, each sodium atom loses 1 electron.
Transition metals
This document discusses the nomenclature and isomers of coordination complexes. It begins by defining ligands and describing systems for naming complexes based on the central metal ion, ligands, and ligand coordination. It then describes three types of isomers that can occur in complexes: structural isomers which have different ligands, linkage isomers which have the same ligands bonded differently, and stereoisomers which have the same ligands in different spatial arrangements. The document goes on to describe crystal field theory and ligand field theory, which explain the splitting of d-orbital energies in an octahedral ligand field based on ligand type. Strong-field ligands create a large energy gap between t2g and eg* orbitals, while weak-field ligands create
Adv chem chapt 9
The molecular orbital model can be used to predict molecular structure, bonding properties, and magnetism. For diatomic molecules, atomic orbitals combine to form molecular orbitals through constructive or destructive interference. Bonding molecular orbitals are lower in energy than antibonding orbitals. Electron configuration and bond order indicate bond strength and stability. Heteronuclear diatomics form molecular orbitals between different atomic orbitals, while delocalized π bonding is seen in resonant structures like benzene.
Metal carbonyls
Metal carbonyls are coordination complexes of transition metals with carbon monoxide ligands. They were first synthesized in 1868 by passing carbon monoxide over platinum. Metal carbonyls typically obey the 18 electron rule and are often diamagnetic. They have applications as catalysts in organic synthesis and in producing pure metals like nickel. Precautions must be taken when using metal carbonyls due to their toxicity.
Naming ionic compounds writing emp formulas
This document discusses how to write empirical formulas for ionic compounds. It explains that ionic compounds are electrically neutral, meaning the positive charge of cations must cancel out the negative charge of anions. It provides an example of determining the charges on potassium (K+) and bromine (Br-) ions and writing the formula for the ionic compound potassium bromide (KBr). The document also gives a mini quiz example of determining the charges on lithium (Li+) and fluorine (F-) ions and writing the formula for the ionic compound lithium fluoride (LiF).
Chemical Bonding and Formula Writing
The document discusses chemical bonding and formula writing. It explains that there are two main types of bonds: ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons between atoms to form ions, while covalent bonds involve the sharing of electrons between atoms. It provides examples of common ionic compounds and their formulas, such as NaCl, as well as explaining how to write formulas for compounds containing polyatomic ions like OH-. Covalent bonding is demonstrated through the example of methane, CH4, where carbon shares electrons with four hydrogen atoms to fill their outer energy levels.
Naming ionic compounds intro
This document discusses the formation of ionic bonds between ions. Ionic bonds are electrostatic attractions between oppositely charged ions. Cations are positively charged ions formed when metals lose electrons. Anions are negatively charged ions formed when nonmetals gain electrons. Ionic bonds generally form between metallic cations and nonmetallic anions. The document provides examples of common cations and anions and identifies metals, nonmetals, and metalloids on the periodic table.
Chemistry - Chp 9 - Chemical Names and Formulas - Notes
1) This document provides information on naming and writing formulas for ionic and molecular compounds. It discusses monatomic and polyatomic ions, and how to determine the charges on cations and anions.
2) Rules are provided for writing formulas and naming ionic compounds, including those containing transition metals. Prefixes are introduced for naming molecular compounds based on the number of atoms present.
3) Examples are worked through for writing formulas from names and names from formulas for different types of compounds. Practice problems are also included for students to try.
2012 ppt unit 2 3 ionic bonding djy r1
Ionic bonding results from the electrostatic attraction between oppositely charged ions. When a metal atom reacts with a nonmetal, the metal typically loses electrons to form a cation while the nonmetal gains electrons to form an anion. These oppositely charged ions are then attracted to each other, forming an ionic bond. Ionic compounds have crystalline structures where the ions are arranged in repeating patterns. Their strong ionic bonds make them brittle with high melting points. Many ionic compounds dissolve in water to form electrolyte solutions where the ions are free to move.
Chemistry - Chp 9 - Chemical Names and Formulas - PowerPoint
This document covers naming and writing formulas for ions, ionic compounds, molecular compounds, acids, and bases. It provides objectives and definitions for each section, examples of naming and writing formulas, and explanations of naming conventions and rules including prefixes, charges, and endings for different types of compounds.
Naming & Writing Ionic Formulas PowerPoint
The document provides information on naming ionic compounds and writing their formulas. It discusses the rules for writing formulas, such as ensuring the total positive and negative charges are equal. It describes the "swoosh and reduce" method for determining formulas, which involves writing the symbols and charges of the elements and crossing the charges to become subscripts on the other element. The document also covers naming cations and anions and how to combine them to name complete ionic compounds.
Back donation
This presentation discusses back-donation of CO ligands to transition metals. Back-donation occurs when electrons from the metal's d-orbitals are donated into the CO ligand's anti-bonding π* orbital, weakening the carbon-oxygen bond. CO can bind to single metals through both sigma and pi bonding. It can also bridge between multiple metal atoms, further weakening the carbon-oxygen bond. Infrared spectroscopy shows that as more metals are involved in back-donation to a CO ligand, its carbon-oxygen stretching frequency decreases due to the bond being more weakened.
Theory of coordination_compounds[1]
The document discusses the theory of coordination compounds. It defines coordination compounds as products of Lewis acid-base reactions where ligands bond to a central metal atom via coordinate covalent bonds. It describes coordination bonds and terminology like ligands, central metal ion, oxidation state, and coordination number. The document also discusses Werner's theory of coordination complexes, valence bond theory, types of ligands, nomenclature of coordination compounds, and limitations of Werner's theory.
transition element
Transition metals have several characteristic properties including high melting points, hardness, and electrical conductivity. They are found in the d-block of the periodic table. The colors of transition metal compounds are determined by the energies absorbed during electron transitions between d orbitals. Spin allowed transitions that don't change the number of paired/unpaired electrons absorb lower energies, appearing as red/yellow colors. Spin forbidden transitions changing this pairing absorb higher energies, appearing as blue/green colors. Coordination number and bond strength also impact the energies absorbed and thus the colors.
Valence Bond Theory
Valence Bond Theory (VBT) explains the nature of metal-ligand bonding in complex compounds. VBT assumes that the central metal atom forms hybrid orbitals with its outer shell orbitals that overlap with ligand orbitals to form sigma bonds. There are two types of hybridization that result in octahedral geometry: d2sp3 for "inner-orbital" complexes with paired electrons, and sp3d2 for "outer-orbital" complexes with unpaired electrons. Examples that form via d2sp3 hybridization, making them inner-orbital complexes, include ferricyanide ([Fe(CN)6]3-), ferrocyanide ([Fe(CN)6]4-
Chapter 6 Notes
Chemical bonds form between atoms to achieve more stable arrangements with lower potential energy. The main types of bonds are ionic, covalent, and metallic. Ionic bonds form between oppositely charged ions, covalent bonds form when atoms share electrons, and metallic bonds result from delocalized electrons being shared among many atoms in a metal. The degree of ionic or covalent character in a bond depends on the electronegativity difference between the atoms. Molecular compounds are made of molecules held together by covalent bonds, while ionic compounds are made of ionic bonds between cations and anions.
Formula writing jhe
This document discusses the steps and rules for writing chemical formulas. It begins by outlining the objectives of learning to write formulas and defining key terms like subscripts and coefficients. It then explains the four main steps to write formulas: 1) Determine element symbols 2) Predict oxidation states 3) Write symbols with charges 4) Balance charges. The rest of the document outlines four main rules for writing subscripts and parentheses correctly based on oxidation states and whether elements appear only once or multiple times in a compound.
Bonding in coordination compound(werners theory)
This document discusses Werner's theory of coordination compounds and bonding in coordination compounds. According to Werner's theory, metal atoms in coordination compounds have both primary and secondary valencies. Primary valencies are ionizable and satisfy the compound's oxidation state, while secondary valencies are non-ionizable and satisfy the compound's coordination number through coordinate covalent bonds with electron pair donors like ligands. The document also discusses Sidgwick's effective atomic number rule and how the valence bond theory explains the geometry, hybridization, and magnetic properties of coordination compounds.
Lecture 9.1 & 9.2- Naming Salts
The document discusses chemical names and formulas of ions. It explains that monatomic cations form when metals lose electrons and gain a positive charge equal to their group number. Monatomic anions form when non-metals gain electrons and have a negative charge equal to the group number minus 8. Polyatomic ions consist of multiple atoms and have an overall positive or negative charge. Transition metals can form multiple ions with different charges that must be specified. Ionic compounds are named by writing the cation first followed by the anion.
Ch 8 Ionic Compounds
1. Ionic compounds form when metals react with nonmetals, resulting in the transfer of electrons and formation of cations and anions bonded ionically.
2. Ionic compounds have high melting points and boiling points due to the strong electrostatic forces between cations and anions in the crystalline lattice structure.
3. Metallic bonds form when metal cations attract delocalized electrons in the "sea of electrons" that moves throughout the entire metallic crystal. This produces metallic properties like conductivity and ductility.
What's hot (20)
Chemistry - Chp 9 - Chemical Names and Formulas - Notes
Chemistry - Chp 9 - Chemical Names and Formulas - Notes
Chemistry - Chp 9 - Chemical Names and Formulas - PowerPoint
Chemistry - Chp 9 - Chemical Names and Formulas - PowerPoint
Viewers also liked
#12 Key
1. The document defines key chemistry terms including: aqueous solutions, solvents, solutes, electrolytes, nonelectrolytes, strong/weak electrolytes, and solvation.
2. Precipitation reactions are defined as reactions where an insoluble product called a precipitate forms. Molecular, complete ionic, spectator ions, and net ionic equations are also defined.
3. Strong acids and bases are listed as well as examples of soluble and insoluble compounds in water. Spectator ions are identified for sample precipitation reactions.
#22 Key
1. An object can possess energy through kinetic energy from its motion or potential energy from its position. Kinetic energy depends on mass and velocity while potential energy depends on position.
2. A kilowatt-hour equals 3.6 million joules. An adult radiates heat at the same rate as a 100-watt bulb, so in 24 hours they radiate about 2.1 kilocalories of heat energy.
3. Chemical reactions that release heat are exothermic, while those that absorb heat are endothermic. Heat change equals internal energy change during constant pressure processes.
#11 Key
The document provides examples of determining empirical and molecular formulas from elemental compositions, balancing chemical equations, calculating theoretical and percent yields of chemical reactions. It includes the following examples:
1) Determining the empirical formulas K3PO4 and Na2SiF6 from elemental compositions.
2) Calculating molecular formulas of H2C2O4 and C4H8O2 given empirical formulas and molar masses.
3) Balancing the equation Fe2O3 + CO → Fe + CO2 and calculating grams of reactants and products.
Metals - Reactivity Series
The document discusses the properties and reactivity of metals. It describes experiments to determine the reactivity series of metals by observing their reactions with water, steam, and dilute acids. Metals react differently in each test based on their positions in the reactivity series, from most reactive to least reactive. The reactivity series allows prediction of other reactions like reduction of metal oxides and decomposition of metal carbonates.
Invasion species
This document discusses invasive species and overexploitation as threats to biodiversity. It defines invasive species as exotic organisms introduced by humans that spread naturally and outcompete native species. Many examples are given of invasive species like eucalyptus in Pakistan, paper mulberry in Islamabad, and zebra mussels in the Great Lakes that have caused environmental and economic harm. Overexploitation is defined as the unsustainable use of natural resources and examples include the commercial whaling industry driving whale populations down and overfishing of bluefin tuna. Control methods proposed include regulating harvesting territories, seasons, and banning inefficient practices.
#10 Key
This document contains 7 problems related to general chemistry. It provides chemical equations to balance, asks students to calculate formula weights and percent compositions of substances, defines key concepts like moles and molar mass, and has multi-step calculations involving moles, molar mass, and grams. The problems cover topics like stoichiometry, formula weights, percent composition, gas laws, and molar conversions.
#32
The document contains 8 questions about general chemistry concepts including:
1) Bond angles in methane and perchlorate based on their tetrahedral electron-domain geometry. NH3's trigonal pyramidal geometry requires specifying one parameter.
2) Determining the number of electron domains and the difference between bonding and nonbonding domains.
3) The characteristic electron-domain geometry of 3, 4, 5, and 6 electron domains around a central atom.
4) The number of electron domains implied by given bond angles.
5) Predicting if molecules are polar or nonpolar based on their structure.
6) The difference between sigma and pi bonds and which is generally stronger.
7) Details about hybrid
#23
The document is a study guide for a general chemistry exam that includes 8 questions covering topics like:
1) Calculating enthalpy changes for chemical reactions using given enthalpy of formation values
2) Calculating standard enthalpy change values (ΔH°) for chemical reactions using reference tables
3) Calculating standard enthalpy of formation for diborane gas using thermochemical data
4) Defining fuel value and comparing the energy provided by fat and carbohydrates
5) Calculating the caloric content of alcohol in wine using given ethanol content and heat of combustion values
6) Explaining why fats are suitable for energy storage and calculating the percentage of calories from fat in a
#21 Key
The document defines key terms related to thermodynamics and calorimetry, including enthalpy, enthalpy of reaction, calorimetry, heat capacity, bomb calorimeter, Hess's law, enthalpy of formation, and standard enthalpy of reaction. It then provides examples of using these concepts to calculate enthalpy changes, heat transferred, and moles or mass of substances involved in various chemical reactions.
Practice test 1 Key
This document contains a practice test for General Chemistry II with 15 multiple choice questions and 5 true/false questions testing concepts related to gases, gas laws, and intermolecular forces. It also includes a bonus section matching chemical names and formulas. The key indicates the correct answers to all questions.
#26 Key
1. The document provides steps to convert between solubility in grams per liter, molar solubility in mol/L, molar concentrations of ions, and Ksp.
2. It gives the definitions of a reducing agent, oxidizing agent, and half-reaction. A reducing agent causes reduction, an oxidizing agent causes oxidation. Half-reactions involve electrons.
3. An anode is where oxidation occurs and a cathode is where reduction occurs in a voltaic cell. Electrons flow from the anode through the external circuit to the cathode.
#23 Key
1. Calculate the enthalpy change for the reaction P4O6 (s) + 2 O2 (g) → P4O10 (s) using given enthalpy values for related reactions.
2. Calculate the standard enthalpy of reaction (ΔH°) for several reactions using standard enthalpies of formation from appendix C.
3. Calculate the standard enthalpy of formation of gaseous diborane (B2H6) using thermochemical cycles and given reaction enthalpies.
4. Explain the meaning of "fuel value" and determine whether 5 g of fat or 9 g of carbohydrates provides more energy as food using their fuel
Viewers also liked (12)
Similar to #31 Key
IONIC-VS-COVALENT-BOND-QUIZ-WITH-ANSWERS.pptx
This document contains 20 multiple choice questions about the types of bonds formed between different elements and how they are formed. It addresses topics like ionic bonds forming between metals and non-metals via electron transfer, covalent bonds forming between non-metals by electron sharing, and the properties of ions and ionic compounds.
week2-d3-Introduction to Bonding.pdf
This document provides an introduction to bonding, including:
- How chemical bonds form via the sharing or transfer of valence electrons between atoms
- Common monoatomic ions like Na+, Ca2+, Cl-, and their charges
- How to write formulas for ionic compounds using charge balance
- Naming conventions for ionic compounds containing metals that form single ions or those with variable charges
- Introduction of polyatomic ions and how to name compounds containing them
7. chemical formulas and naming compnds [Autosaved].pptx
This document provides information on naming and writing formulas for different types of compounds. It discusses binary ionic compounds and how their names and formulas are determined based on the ions present. It also covers binary covalent compounds and how prefixes are used in their names to indicate the ratio of elements. Finally, it mentions polyatomic ions and how their names are used in ionic compound names and formulas.
Lecture 7.1- Ions
Magnesium and chlorine have valence electrons in their highest energy level. Magnesium forms Mg2+ ions by losing two valence electrons to attain the electron configuration of neon. Chlorine forms Cl- ions by gaining one electron to attain the electron configuration of argon. Ions are formed by atoms gaining or losing electrons to achieve a noble gas configuration.
Transition metals
This document discusses the nomenclature and isomers of coordination complexes. It begins by defining ligands and describing systems for naming complexes based on the central metal ion, ligands, and ligand coordination. It then describes three types of isomers that can occur in complexes: structural isomers which have different ligands, linkage isomers which have the same ligands bonded differently, and stereoisomers which have the same ligands in different spatial arrangements. The remainder of the document outlines crystal field theory and ligand field theory to explain electron configurations, orbital splitting patterns, and spectroscopic properties of transition metal complexes based on the geometry and ligand type.
#30
The document contains questions about trends in the periodic table including ion size, isoelectronic series, electron configuration of ions, ionization energy, electron affinity, reactivity of alkali metals, balancing equations for alkali metal oxide reactions, and properties of noble gases. It asks about trends in properties including ionic radius, ionization energy, metallic character, and molecular structure as one moves through the periodic table.
Chembond
Chemical bonds form between atoms in different ways depending on electron configuration. Ionic bonds form when electrons are transferred between metals and nonmetals, creating positively and negatively charged ions. Covalent bonds form when electrons are shared between nonmetals. Metallic bonds form between metal atoms through a "sea of electrons" that holds the atoms together.
Chemistry - Chp 7 - Ionic And Metallic Bonding - PowerPoint
This document summarizes key concepts from Chapter 7 on ionic and metallic bonding. It discusses how ions form as atoms gain or lose electrons to achieve stable noble gas electron configurations. Cations are positive ions that form when metals lose electrons, while anions are negative ions that form when nonmetals gain electrons. Ionic compounds consist of cations and anions bonded through electrostatic attraction. Metals form metallic bonds where valence electrons are delocalized and allow the structure to be malleable. Alloys are mixtures of metals that can have superior properties to pure metals.
Chemistry - Chp 7 - Ionic and Metallic Bonding - PowerPoint
This document summarizes key concepts from Chapter 7 on ionic and metallic bonding. It discusses how ions form as atoms gain or lose electrons to achieve stable noble gas electron configurations. Cations are positive ions that form when metals lose electrons, while anions are negative ions that form when nonmetals gain electrons. Ionic compounds consist of cations and anions bonded through electrostatic attraction. Their properties include being crystalline solids with high melting points that conduct electricity when molten or dissolved. Metallic bonding is described as a "sea of electrons" where metal atoms are positively charged cations floating in a sea of delocalized electrons that are free to move, explaining metals' malleability and conductivity. Alloys are mixtures of
C4 And C5 Multiple Choice Revision Quiz
This document contains a 20 question multiple choice quiz about concepts related to atomic structure, bonding, and compounds. The questions cover topics like ionic bonding, metallic bonding, properties of ionic and covalent compounds, electron configuration, and electrolysis.
Ionic Bonds - Chapter 7
Ionic bonding occurs when atoms transfer electrons to form ions with opposite charges that are attracted via electrostatic forces. Metals form cations by losing electrons to achieve stable electron configurations like noble gases, while nonmetals form anions by gaining electrons. This transfer of electrons allows the formation of ionic compounds with crystalline structures where ion attractions are maximized and repulsions minimized. Properties of ionic compounds include high melting points, solubility in water, defined crystal structures, and the ability to conduct electricity when molten. Metallic bonding also involves cations but is characterized by delocalized valence electrons that form a "sea" allowing metals to conduct electricity and be malleable and ductile.
Chembond
Chemical bonds form between atoms in order to fill their outer electron shells. Ionic bonds form when electrons are transferred between metals and nonmetals, creating positively and negatively charged ions. Covalent bonds form when electrons are shared between nonmetals. Metallic bonds form between metal atoms through a "sea of electrons".
Ch 8 ionic compounds
This document discusses ionic compounds and the formation of ionic bonds. Ions form when atoms gain or lose electrons to achieve stable electron configurations like noble gases. Ionic bonds occur between oppositely charged ions and result in crystalline solids with high melting points. The document explains how to name ionic compounds based on the cation and anion present and write chemical formulas from compound names. It also briefly discusses metallic bonding and the properties of metals and alloys.
2 Atomic Structure
1. The document discusses atomic structure and properties of subatomic particles like electrons, protons, and neutrons.
2. It also covers terminology used in atomic structure like atomic number, mass number, and isotopes. Methods for calculating relative atomic mass and isotope percentages are presented.
3. The formation of ionic and covalent bonds is explained through examples of ionic compounds like NaCl and covalent compounds like H2O and CH4. Dot and cross diagrams are used to represent bonding.
Ch 8 ionic compounds
Ionic compounds form when a metal transfers electrons to a nonmetal, resulting in oppositely charged ions that are attracted in a crystal lattice structure. The ratio of cations to anions must result in an overall neutral charge. Ionic compounds are named by writing the cation first followed by the anion. Polyatomic ions are named based on the nonmetal and number of oxygen atoms. Metallic bonds form a "sea of electrons" that are delocalized and allow metal atoms to strongly attract each other.
Chapter 6: Bonding Basics
If you have been absent, view the class PowerPoint for Chapter 6. You should be able to view it directly on FB.
9380673.ppt
The document discusses coordination compounds and transition metals. It begins by defining transition metals as having incompletely filled d subshells. It then discusses the electronic structures of zinc and zinc ions to determine if zinc is a transition metal. The rest of the document covers topics like common ligands, naming coordination compounds, the various structures they can form, and crystal field theory which explains their bonding. It provides examples of ligand substitution reactions and applications of coordination compounds.
Ikatan kimIa Arum
1. The document discusses different types of chemical bonding including ionic bonding, covalent bonding, and metallic bonding.
2. Ionic bonding occurs when atoms transfer electrons to form oppositely charged ions that are then attracted via electrostatic forces, such as in sodium chloride (NaCl).
3. Covalent bonding involves the sharing of electron pairs between atoms to form molecular compounds, like in water (H2O) and methane (CH4).
What are groups and periods?
This document is a lesson on groups and periods in the periodic table. It defines periods as rows from left to right, and groups as columns from top to bottom. Elements in the same group have similar properties because they have the same number of valence electrons. The lesson discusses special metal and nonmetal groups, including alkali metals, alkaline earth metals, transition metals, halogens, and noble gases. It provides examples of elements in each group and their properties. The document concludes with practice questions for students.
Molecular orbital theory
Molecular orbital theory (MOT) is an alternative model to valence bond theory that explains how atomic orbitals from different atoms combine to form molecular orbitals. The linear combination of atomic orbitals (LCAO) method considers the probability of finding electrons in atomic orbitals from different atoms. According to the LCAO method, molecular orbitals are formed from constructive and destructive interference of atomic orbitals. MOT can be used to explain bonding in homonuclear diatomic molecules like N2 and O2, heteronuclear diatomic molecules like CO and NO, and polyatomic molecules like CO2 and SF6. It can also describe bonding in octahedral transition metal complexes like hexaaquoferrate(II) ion
Similar to #31 Key (20)
7. chemical formulas and naming compnds [Autosaved].pptx
7. chemical formulas and naming compnds [Autosaved].pptx
Chemistry - Chp 7 - Ionic And Metallic Bonding - PowerPoint
Chemistry - Chp 7 - Ionic And Metallic Bonding - PowerPoint
Chemistry - Chp 7 - Ionic and Metallic Bonding - PowerPoint
Chemistry - Chp 7 - Ionic and Metallic Bonding - PowerPoint
More from Lamar1411_SI
#29 Key
The document discusses key concepts in general chemistry including:
1) Effective nuclear charge increases left to right and down a period, lowering electron energy due to stronger electrostatic attraction but increasing it due to greater electron-electron repulsion.
2) Magnesium has an electron configuration of [Ne]3s2, with 2 valence electrons and an effective nuclear charge of +2.
3) Atomic radius generally decreases left to right and increases down a period, while ionic radius follows the opposite trends due to shielding effects.
#29
The document discusses concepts in general chemistry including effective nuclear charge, electron configuration, atomic and ionic sizes, and ionization energies. It defines effective nuclear charge as the charge felt by an electron from the nucleus and how it decreases across a period as nuclear charge is shared by more electrons. It asks about the electron configuration, valence electrons, and energy level of magnesium as well as how its effective nuclear charge differs from carbon. Finally, it asks about writing equations for specific ionization energies of tin and titanium.
#27 Key
1. The maximum number of electrons that can occupy each subshell is: 3p = 6, 5d = 10, 2s = 2, 4f = 14.
2. The maximum number of electrons for the given quantum numbers are: a) n = 2, ms = -1/2 is 3 electrons, b) n = 5, l = 3 is 14 electrons, c) n = 4, l = 3, ml = -3 is 2 electrons, d) n = 4, l = 1, ml = 1 is 2 electrons.
3. The electromagnetic spectrum ranges from high energy gamma rays and x-rays to lower energy radio waves.
#27
This document contains 6 questions about general chemistry concepts including:
1) The maximum number of electrons that can occupy atomic subshells and quantum numbers
2) Writing condensed electron configurations and determining unpaired electrons for various atoms
3) Calculating frequency and wavelength from each other
4) Labeling statements about electromagnetic radiation as true or false
5) Calculating and comparing photon energies at different wavelengths
6) Drawing the electromagnetic spectrum.
#26 Key
The document provides definitions and explanations of key concepts in quantum mechanics and atomic structure, including:
- The formula for the energy of a photon
- Definitions of wave function, probability density, electron density, orbitals, electron shells, subshells, degeneracy, Pauli exclusion principle, electron configuration, and Hund's rule
- Tables listing the allowed values of the principal, angular momentum, magnetic, and spin quantum numbers
- Examples of specifying orbitals using n, l, and ml values
- Examples of writing electron configurations in both long and short hand notation for various elements
#26
The document contains questions about fundamental concepts in general chemistry including:
1) The formula for the energy of a photon.
2) The ground and excited states of a photon.
3) Definitions of key terms like wave function, probability density, orbitals, electron shells, subshells, degeneracy, Pauli exclusion principle, electron configuration, and Hund's rule.
4) Tables to fill in about quantum numbers and maximum possible electrons in different orbitals.
5) Questions about the values of quantum numbers n, l, and ml for different orbitals.
6) Determining electron configurations using both longhand and shorthand notation for various elements.
#22
1. An object can possess energy in two ways: potential energy due to position or configuration, and kinetic energy due to motion. Potential energy is stored energy due to factors like height or chemical composition, while kinetic energy is energy due to an object's motion.
2. A kilowatt-hour is equal to 3,600,000 joules. An adult radiates heat at about 100 watts, so in 24 hours would radiate 24,000 kcal of energy to the surroundings.
3. For the chemical reaction of silver ions and chloride ions forming silver chloride: a) the heat of reaction for 0.200 mol of AgCl forming is -13.12 kJ; b) the
#21
This document defines key terms in chemistry including enthalpy, enthalpy of reaction, calorimetry, heat capacity, bomb calorimeter, Hess's Law, enthalpy of formation, and standard enthalpy change. It then provides several practice problems applying these concepts, such as writing thermochemical equations, calculating enthalpy changes, and determining the heat required to change temperatures.
#20 Key
This document defines key terms in thermodynamics and chemistry, including:
- Kinetic energy, potential energy, joules, calories, work, heat, internal energy, the first law of thermodynamics, endothermic and exothermic processes, and state functions.
- It provides examples of calculating kinetic energy, changes in energy for objects moving in different scenarios, and determining whether processes are endothermic or exothermic based on changes in heat and work.
- Key concepts covered include the relationships between motion, position, and potential/kinetic energy, and how the first law of thermodynamics relates the change in internal energy of a system to heat and work.
#20
The document defines key terms in thermodynamics and chemistry including thermodynamics, kinetic energy, potential energy, joules, calories, work, energy, heat, internal energy of a system, the first law of thermodynamics, endothermic process, and exothermic process. It then provides examples to calculate kinetic energy, convert between joules and calories, and determine if processes are endothermic or exothermic by calculating the change in energy and accounting for heat and work.
#19 key
1. 18.9 grams of K2Cr2O7 are present in 50.0 mL of 0.360 M K2Cr2O7 solution.
2. The molarity of the (NH4)2SO4 solution formed from dissolving 4.28 grams of (NH4)2SO4 in 300 mL of water is 0.108 M.
3. 58.7 mL of 0.240M CuSO4 solution contains 2.25 grams of CuSO4 solute.
#19
This document contains 10 chemistry problems involving calculations of molarity, moles, and mass in solutions. It asks the learner to calculate the grams of solute in a given volume of a solution, the molarity of a solution made by dissolving a mass of solute, the volume of a solution containing a given mass of solute, which of several solutions contains the highest concentration of a particular ion, which of two solutions contains the greater number of moles of an ion, the concentrations of individual ions or molecules in various solutions, the mass of one substance needed to precipitate another from a solution, the volume of one acid needed to neutralize a base solution, the volume of acid needed to neutralize a mass of
#18 Key
1. When ions dissolve in water, they are surrounded by water molecules (hydrated).
2. Precipitation will occur when Na2CO3 and AgNO3 or FeSO4 and Pb(NO3)2 are mixed, but not NaNO3 and NiSO4.
3. Acids donate H+ ions, bases accept H+ ions. A monoprotic acid has one ionizable H, a diprotic acid has two. Strong acids fully ionize, weak acids only partially ionize.
#18
1. When an ionic substance dissolves in water, the ions become surrounded by and separated from each other by water molecules.
2. Precipitation will occur when solutions of Na2CO3 and AgNO3 are mixed, forming the insoluble Ag2CO3. Precipitation will also occur when FeSO4 and Pb(NO3)2 are mixed, forming the insoluble PbSO4. No precipitation will occur between NaNO3 and NiSO4.
3. Strong acids and bases completely dissociate in water, while weak acids and bases only partially dissociate. A monoprotic acid contains one ionizable hydrogen per molecule
#17 Key
This document provides sample chemistry problems and questions related to topics like:
- Writing balanced equations for neutralization reactions
- Defining terms like oxidation, reduction, concentration, and indicators
- Determining oxidation states of elements in compounds
- Identifying redox, precipitation, and acid-base reactions
- Performing calculations involving molarity, moles, and concentration
The problems cover a wide range of general chemistry concepts.
#17
The document contains questions about chemistry concepts including:
1) Writing balanced molecular and net ionic equations for neutralization reactions.
2) Defining oxidation, reduction, concentration, molarity, and indicators.
3) Calculating molarity, moles, and volumes for various chemistry problems involving solutions.
#13 Key
This document provides sample problems and questions for a general chemistry exam. It includes sample balanced equations for neutralization reactions and definitions of key terms like oxidation, reduction, concentration, and indicators. It also asks students to determine oxidation states, identify redox reactions, calculate molarity and amounts of substances in solutions, and determine concentrations of ions after mixing solutions.
#13
The document contains questions about chemistry concepts including:
1) Writing balanced molecular and net ionic equations for neutralization reactions between acids and bases.
2) Defining oxidation, reduction, concentration, molarity, and indicators.
3) Asking if oxidation can occur without accompanying reduction.
4) Calculating oxidation numbers of elements in compounds.
5) Identifying redox, precipitation, and acid-base reactions based on balanced equations.
6) Calculating molarity, moles of solute, and volumes of solutions.
Si #12
This document defines key chemistry terms and concepts related to aqueous solutions, electrolytes, acids and bases, precipitation reactions, and solubility. It asks the student to identify strong acids and bases, ions present when various substances dissolve in water, solubility of compounds, precipitation reactions, and classification of solutions as nonelectrolytes, weak electrolytes, or strong electrolytes. The document provides a review of general chemistry topics through multiple choice and short answer questions.
#31 Key
- 1. KEY GENERAL CHEMISTRY-I (1411) S.I. # 31 1. Metal characteristic increases _______B_______ a group. a. up b. down c. across 2. Metal character decreases _______C_______ a period. a. up b. down c. across 3. Metals have ______B________ ionization energies. a. high b. low c. zero d. positive e. negative 4. Group 1A metals for _____M+____ Ions and group 2A form ____M2+_____ions. 5. Most metal oxides are _____B______ a. acidic b. basic c. neutral 6. Most non metal oxides are: ____A______ a. acidic b. basic c. neutral 7. When nonmetal oxides react with water what is formed? An Acid 8. Complete and balance the following equations. P4O10(s) + H2O(l) 4H3PO4(aq) CO2(g) + H2O(l) H2CO3(aq) 9. What are some properties of metalloids? Metalloids have properties that are intermediate between metals and nonmetals. Metalloids have found fame in the semiconductor industry. 10. Using only the periodic table, arrange the following atoms in increasing radius (a) Cs, K, Rb K < Rb < Cs__________________ (b) In, Te, Sn Te < Sn < In_________________ (c) P, Cl, Sr Cl < P < Sr___________________ 11. What is the octet rule? Atoms will gain, lose or share electrons to achieve the nearest noble gas electron configuration. Except for H and He, this corresponds to eight electrons in the valence shell, thus the term octet rule 12. How many electrons must a sulfur atom gain to achieve an octet in its valence shell? S: [Ne]3s23p4 A sulfur atom has six valence electrons, so it must gain two electrons to achieve an octet. 13. If an atom has the electron configuration 1s22s22p3, how many electrons must it gain to achieve an octet? 1s22s22p3 = [He]2s22p3 The atom (N) has five valence electrons and must gain three electrons to achieve an octet
- 2. KEY 14. a. Write the electron configuration for the element scandium, Sc. Sc: 1s22s22p63s23p64s23d1 = [Ar]4s23d1 b. How many valence electrons does this atom possess? Scandium has three (3) valence electrons. c. What distinguishes these valence electrons from the other electrons in the atom? These valence electrons are available for chemical bonding, while the core electrons do not participate in bonding. 15. Use Lewis symbols to represent the reaction that occurs between Mg and Br atoms. . . . . . . . . . . Mg +. Br . +. Br . . Mg2+ + 2[ .Br. ] - . . . . . . . 16. What ionic compound is expected to form from combination of the following pairs of elements? a. barium and oxygen BaO___________________ b. rubidium and iodine RbI ___________________ c. lithium and sulfur Li2S ___________________ d. bromine and magnesiumMgBr2___________________ 17. Which of these elements is unlikely to form covalent bonds: S, H, K, Ar, Si? Explain your choices K and Ar. K is an active metal with one valence electron . It is most likely to achieve an octet by losing this single electron and to participate in ionic bonding. Ar has a stable octet of valence electrons; it is not likely to form chemical bonds of any type. 18. a. What is the most electronegative element in group 6A? O_____ b. Which is most electronegative out of Al, Si, P? P_____ c. Which is most electronegative out of Ga, P, Cl, Na? Cl_____ d. The element most likely to form an ionic compound with Ba out of K, C, Zn, F F 19. Arrange the bonds in each of the following sets in order of increasing polarity: a. C--F, O—F, Be—F O—F < C—F < Be—F b. O—Cl, S—Br, C—P S—Br < C—P < O—Cl c. C—S, B—F , N—O C—S < N—O < B—F 20. Does the octet rule apply to ionic as well as covalent compounds? Explain. The octet rule applies to the individual ions in an ionic compound. A cation loses electrons and the anion gains electrons to achieve an octet. For example: in MgCl2, Mg loses 2 e- to become Mg2+ with the electron configuration of [Ne]. Each Cl atom gains one electron to form Cl- with the electron configuration of [Ar]. 21. Draw the Lewis structures for each of the following molecules or ions. Which do not obey the octet rule?
- 3. KEY a. CO2 b. IO3- c. BH3 d. BF4- e. XeF2 A and D