1. The document is a periodicity chemistry worksheet containing questions about periodic trends, atomic structure, ionization energy, electronegativity, and other periodic properties.
2. Key periodic trends addressed include ionization energy increasing across a period and decreasing down a group, atomic radius decreasing across a period and increasing down a group, and electronegativity generally increasing moving up and to the right on the periodic table.
3. Factors that influence periodic trends are discussed, such as the shielding effect increasing with greater nuclear charge, which causes atomic radius to decrease and ionization energy to increase moving across a period.
This document is a periodic trends worksheet that contains questions to test understanding of trends in atomic radius, ionization energy, and electronegativity across the periodic table. It asks students to rank elements by atomic radius and electronegativity. It also asks students to identify whether atomic radius, ionization energy, and electronegativity increase or decrease moving left to right or top to bottom on the periodic table. Finally, it asks students to identify which atom in pairs has the larger/smaller atomic radius, ionization energy, or electronegativity.
At the end of this chapter you should be able to sketch the periodic table showing the groups and periods; identify the metals, metalloids and non-metals in the periodic table. Identify the representative elements, the transition elements, the transition metals, the lanthanides and actinides in the periodic table. Also, give the electron configuration of cations and anions; determine the trends in the physical properties of elements in a group; describe and explain the trends in atomic properties in the periodic table; compare the properties of families and elements; predict the properties of individual elements based on their position in the periodic table; and perform exercises and collaborative work with peers.
The document discusses trends in the properties of elements in the periodic table. It describes how elements in the same group generally have similar chemical properties, though properties are not identical. It discusses how Mendeleev and Meyer independently organized the elements into the first periodic tables based on recurring trends in properties. Elements in the inner transition metals block were later added.
This document provides an overview of chemical bonding. It discusses valence electrons and how to determine the number for each element using the periodic table. It describes the different types of chemical bonds - ionic bonds formed by the transfer of electrons between atoms, and covalent bonds formed by the sharing of electron pairs. It also discusses concepts like electronegativity and how it relates to bond type and strength.
Interactive textbook ch. 13 chemical bondingtiffanysci
Ionic bonds form when valence electrons are transferred from metal atoms to nonmetal atoms, resulting in positively charged metal ions and negatively charged nonmetal ions that are attracted to each other. Metal atoms easily lose electrons to achieve stable full outer energy levels, while nonmetal atoms gain electrons for the same reason. The ions associate in repeating three-dimensional crystal lattices to form solid ionic compounds that are brittle with high melting points and often dissolve in water.
A nearly-comprehensive list of vocabulary terms needed for introductory chemistry in grade 9 science, including a variety of source websites for reference.
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.
This document is a periodic trends worksheet that contains questions to test understanding of trends in atomic radius, ionization energy, and electronegativity across the periodic table. It asks students to rank elements by atomic radius and electronegativity. It also asks students to identify whether atomic radius, ionization energy, and electronegativity increase or decrease moving left to right or top to bottom on the periodic table. Finally, it asks students to identify which atom in pairs has the larger/smaller atomic radius, ionization energy, or electronegativity.
At the end of this chapter you should be able to sketch the periodic table showing the groups and periods; identify the metals, metalloids and non-metals in the periodic table. Identify the representative elements, the transition elements, the transition metals, the lanthanides and actinides in the periodic table. Also, give the electron configuration of cations and anions; determine the trends in the physical properties of elements in a group; describe and explain the trends in atomic properties in the periodic table; compare the properties of families and elements; predict the properties of individual elements based on their position in the periodic table; and perform exercises and collaborative work with peers.
The document discusses trends in the properties of elements in the periodic table. It describes how elements in the same group generally have similar chemical properties, though properties are not identical. It discusses how Mendeleev and Meyer independently organized the elements into the first periodic tables based on recurring trends in properties. Elements in the inner transition metals block were later added.
This document provides an overview of chemical bonding. It discusses valence electrons and how to determine the number for each element using the periodic table. It describes the different types of chemical bonds - ionic bonds formed by the transfer of electrons between atoms, and covalent bonds formed by the sharing of electron pairs. It also discusses concepts like electronegativity and how it relates to bond type and strength.
Interactive textbook ch. 13 chemical bondingtiffanysci
Ionic bonds form when valence electrons are transferred from metal atoms to nonmetal atoms, resulting in positively charged metal ions and negatively charged nonmetal ions that are attracted to each other. Metal atoms easily lose electrons to achieve stable full outer energy levels, while nonmetal atoms gain electrons for the same reason. The ions associate in repeating three-dimensional crystal lattices to form solid ionic compounds that are brittle with high melting points and often dissolve in water.
A nearly-comprehensive list of vocabulary terms needed for introductory chemistry in grade 9 science, including a variety of source websites for reference.
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.
The document discusses the development and key features of the modern periodic table of elements. It describes how early scientists like Lavoisier, Dobereiner, Newlands, Meyer, and Mendeleev organized the known elements and recognized patterns in their properties. Moseley's work in the early 1900s led to developing the modern periodic law based on atomic number. The periodic table is now divided into periods and groups with characteristic trends in properties like atomic radius, ionization energy, and electronegativity that can be explained by electron configuration.
The document discusses the periodic table and periodic trends in properties of elements. It defines key periodic table terms like groups, periods, atomic radius, ionization energy, and provides examples of how properties change within periods and groups. It also discusses Mendeleev's original periodic table and the modern periodic law.
The document discusses different types of chemical bonds including ionic bonds, covalent bonds, and polar covalent bonds. It describes how ionic bonds form between a metal and nonmetal when electrons are transferred, covalent bonds form through shared electron pairs, and polar covalent bonds result in an unequal sharing of electrons. The document also covers bond energies, lattice energies in ionic compounds, electronegativity, and molecular polarity.
10/13 Review: What is electronegativity and ionization energy?mrheffner
The document summarizes a chemistry lesson on periodic trends. It discusses the trends in atomic radius, ionization energy, and electronegativity across the periodic table. For atomic radius, it decreases from left to right and increases from top to bottom. Ionization energy increases from left to right and decreases from top to bottom. Electronegativity also increases from left to right and decreases from top to bottom. Examples are provided to demonstrate how to apply these trends to determine which atom has a larger/smaller value of the property. An exit slip quiz and homework assignment are assigned.
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.
test bank Organic Chemistry, 7e Marc Loudon, Jim Parise test bank.pdfNailBasko
This document appears to be the beginning of a chemistry textbook chapter, containing questions and answers about various chemistry concepts. It includes 25 multiple choice or short answer questions covering topics like molecular orbitals, Lewis structures, formal charges, bond types, electronegativity, and molecular geometry. The questions progress from basic concepts like electron configurations to more advanced topics involving resonance and hybridization.
This document provides information about the periodic table and periodic properties of elements. It discusses the contributions of Dobereiner, Newlands and Mendeleev towards developing the periodic table. It explains that elements are arranged in the modern periodic table according to increasing atomic number. Properties like atomic radius, ionization energy and electronegativity repeat periodically with this arrangement. The document also answers multiple choice and short answer questions about periodic trends, periodic law, and classification of elements in periods and groups.
The document provides information about chemical bonds including ionic bonds, covalent bonds, and bond energies. It defines ionic and covalent bonding, discusses factors that determine lattice energy of ionic compounds, introduces electronegativity and bond polarity. It also covers Lewis structures, resonance structures, and exceptions to the octet rule. Bond enthalpies, which measure bond strength, are discussed along with average bond enthalpies from bond dissociation data.
Atomos, Teorías Atómicas, Teoría Cuántica, Masa Atómica e Isotoposfernandogc
The document provides information about atoms and their subatomic particles. It defines the atom as the smallest particle of an element that retains its characteristics. It describes atoms as consisting of protons, neutrons, and electrons. The number of protons defines the element and its atomic number, while the total of protons and neutrons is the mass number. Isotopes are atoms of the same element with different numbers of neutrons. Electron configuration diagrams show how electrons are arranged in an atom's energy levels.
The document discusses periodic properties of elements including atomic radii, ionization energy, electron affinity, ionic radii, and electronegativity. It provides information on trends in these properties across the periodic table and examples demonstrating how to order elements based on each property. Atomic radii decrease and ionization energy increases moving from left to right within a period due to increased nuclear charge. Electron affinity and ionic radii values become more negative in the same trend.
The document discusses the structure of atoms and includes questions and answers about:
- The discovery of electrons, protons, and neutrons by different scientists.
- How anode rays are formed from gas in a discharge tube.
- Evidence from Rutherford's gold foil experiment that atomic mass is concentrated at the center in a nucleus.
- How Rutherford showed atomic nuclei are positively charged using alpha particle bombardment.
- Particles that determine atomic mass.
- Differences between classical and quantum theories of radiation and angular momentum quantization.
The chemistry objective midterm exam will cover chapters 1-7 and include 50 multiple choice questions. Students should review key concepts from each chapter including the scientific method, density calculations, significant figures, atomic structure, the periodic table, chemical bonding and reactions. They should understand periodic trends in atomic properties and be prepared to classify matter, write formulas and solve stoichiometry problems.
This document discusses key concepts in atomic structure and bonding:
1. It describes John Dalton's atomic theory and the discoveries of J.J. Thomson, Ernest Rutherford, and other scientists that led to developing the modern atomic model.
2. Key aspects of the modern atomic model are explained, including that atoms are composed of protons, neutrons, and electrons, with the nucleus at the center containing protons and neutrons.
3. The document also introduces concepts such as isotopes, ions, and how elements are arranged on the periodic table based on their atomic structure.
This document outlines the key concepts to be covered in a Year 11 100 Science course on aspects of acids and bases, including atomic structure, properties of acids and bases, rates of reaction and particle theory, uses of acids and bases, and restrictions on the acids and bases included in the course. Students will study electron configuration, ionic bonding, naming ionic compounds, properties of acids and bases such as releasing hydrogen ions in water and reacting to form salts, and the rates of reactions and particle theory explanations. Assessment will include selected aspects of acids and bases such as atomic structure, properties, uses, and rates of reaction.
This document discusses the molecular structure of atoms. It begins by explaining that all substances are composed of elements, which are the basic substances, and that elements are made up of very small particles called atoms. It then describes the basic structure of an atom, including that atoms have a small, positively charged nucleus surrounded by negatively charged electrons. The number of protons determines the element and equals the number of electrons. Atoms can gain, lose, or share electrons to attain stable electron configurations, forming ions or molecules.
- Intramolecular forces are attractions between atoms within the same molecule, while intermolecular forces are attractions between different molecules.
- Covalent bonds form when atoms share electrons, and can be polar or non-polar depending on differences in electronegativity between the bonded atoms. Polar covalent bonds have partial charges while non-polar covalent bonds share electrons equally.
- The polarity of a molecule depends on differences in electronegativity between its bonded atoms, with larger differences resulting in more polar bonds and molecules. Water is an example of a polar molecule due to differences in electronegativity between oxygen and hydrogen.
1. Ionic compounds form when a metal reacts with a non-metal, resulting in positively charged metal ions and negatively charged non-metal ions that bond together in a crystalline lattice structure.
2. When ionic compounds dissolve in water or melt, the ions become free to move and conduct electricity. During electrolysis, positively charged metal ions move to the cathode and negatively charged non-metal ions move to the anode.
3. Common ionic compounds include sodium chloride, formed from sodium and chlorine ions, and copper chloride, used in electrolysis to extract copper metal from its ionic form.
This document outlines key concepts about atomic structure including:
1) The structure of atoms consists of a nucleus containing protons and neutrons surrounded by electrons. Protons are positively charged, neutrons have no charge, and electrons are negatively charged.
2) Atomic number refers to the number of protons in an atom. Mass number refers to the total number of protons and neutrons. Atomic symbols represent these numbers.
3) Electrons surround the nucleus in fixed shells. The number of electrons in each shell is limited. The arrangement of electrons is known as the electronic structure or electron configuration.
This document discusses atomic structure and electron configuration. It contains the following key points:
1. Atoms are made up of protons, neutrons, and electrons, with protons and neutrons located in the nucleus and electrons orbiting the nucleus.
2. Electrons can occupy different energy levels characterized by quantum numbers like principal quantum number and azimuthal quantum number. This determines the shape of electron orbitals.
3. The electron configuration of an atom shows the distribution of electrons among these orbitals, following the Pauli exclusion principle. Valence electrons are in the outermost shell and influence chemical bonding and properties.
4. Elements are arranged in the periodic table based on their atomic structure, including number
This presentation includes basic of PCOS their pathology and treatment and also Ayurveda correlation of PCOS and Ayurvedic line of treatment mentioned in classics.
The document discusses the development and key features of the modern periodic table of elements. It describes how early scientists like Lavoisier, Dobereiner, Newlands, Meyer, and Mendeleev organized the known elements and recognized patterns in their properties. Moseley's work in the early 1900s led to developing the modern periodic law based on atomic number. The periodic table is now divided into periods and groups with characteristic trends in properties like atomic radius, ionization energy, and electronegativity that can be explained by electron configuration.
The document discusses the periodic table and periodic trends in properties of elements. It defines key periodic table terms like groups, periods, atomic radius, ionization energy, and provides examples of how properties change within periods and groups. It also discusses Mendeleev's original periodic table and the modern periodic law.
The document discusses different types of chemical bonds including ionic bonds, covalent bonds, and polar covalent bonds. It describes how ionic bonds form between a metal and nonmetal when electrons are transferred, covalent bonds form through shared electron pairs, and polar covalent bonds result in an unequal sharing of electrons. The document also covers bond energies, lattice energies in ionic compounds, electronegativity, and molecular polarity.
10/13 Review: What is electronegativity and ionization energy?mrheffner
The document summarizes a chemistry lesson on periodic trends. It discusses the trends in atomic radius, ionization energy, and electronegativity across the periodic table. For atomic radius, it decreases from left to right and increases from top to bottom. Ionization energy increases from left to right and decreases from top to bottom. Electronegativity also increases from left to right and decreases from top to bottom. Examples are provided to demonstrate how to apply these trends to determine which atom has a larger/smaller value of the property. An exit slip quiz and homework assignment are assigned.
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.
test bank Organic Chemistry, 7e Marc Loudon, Jim Parise test bank.pdfNailBasko
This document appears to be the beginning of a chemistry textbook chapter, containing questions and answers about various chemistry concepts. It includes 25 multiple choice or short answer questions covering topics like molecular orbitals, Lewis structures, formal charges, bond types, electronegativity, and molecular geometry. The questions progress from basic concepts like electron configurations to more advanced topics involving resonance and hybridization.
This document provides information about the periodic table and periodic properties of elements. It discusses the contributions of Dobereiner, Newlands and Mendeleev towards developing the periodic table. It explains that elements are arranged in the modern periodic table according to increasing atomic number. Properties like atomic radius, ionization energy and electronegativity repeat periodically with this arrangement. The document also answers multiple choice and short answer questions about periodic trends, periodic law, and classification of elements in periods and groups.
The document provides information about chemical bonds including ionic bonds, covalent bonds, and bond energies. It defines ionic and covalent bonding, discusses factors that determine lattice energy of ionic compounds, introduces electronegativity and bond polarity. It also covers Lewis structures, resonance structures, and exceptions to the octet rule. Bond enthalpies, which measure bond strength, are discussed along with average bond enthalpies from bond dissociation data.
Atomos, Teorías Atómicas, Teoría Cuántica, Masa Atómica e Isotoposfernandogc
The document provides information about atoms and their subatomic particles. It defines the atom as the smallest particle of an element that retains its characteristics. It describes atoms as consisting of protons, neutrons, and electrons. The number of protons defines the element and its atomic number, while the total of protons and neutrons is the mass number. Isotopes are atoms of the same element with different numbers of neutrons. Electron configuration diagrams show how electrons are arranged in an atom's energy levels.
The document discusses periodic properties of elements including atomic radii, ionization energy, electron affinity, ionic radii, and electronegativity. It provides information on trends in these properties across the periodic table and examples demonstrating how to order elements based on each property. Atomic radii decrease and ionization energy increases moving from left to right within a period due to increased nuclear charge. Electron affinity and ionic radii values become more negative in the same trend.
The document discusses the structure of atoms and includes questions and answers about:
- The discovery of electrons, protons, and neutrons by different scientists.
- How anode rays are formed from gas in a discharge tube.
- Evidence from Rutherford's gold foil experiment that atomic mass is concentrated at the center in a nucleus.
- How Rutherford showed atomic nuclei are positively charged using alpha particle bombardment.
- Particles that determine atomic mass.
- Differences between classical and quantum theories of radiation and angular momentum quantization.
The chemistry objective midterm exam will cover chapters 1-7 and include 50 multiple choice questions. Students should review key concepts from each chapter including the scientific method, density calculations, significant figures, atomic structure, the periodic table, chemical bonding and reactions. They should understand periodic trends in atomic properties and be prepared to classify matter, write formulas and solve stoichiometry problems.
This document discusses key concepts in atomic structure and bonding:
1. It describes John Dalton's atomic theory and the discoveries of J.J. Thomson, Ernest Rutherford, and other scientists that led to developing the modern atomic model.
2. Key aspects of the modern atomic model are explained, including that atoms are composed of protons, neutrons, and electrons, with the nucleus at the center containing protons and neutrons.
3. The document also introduces concepts such as isotopes, ions, and how elements are arranged on the periodic table based on their atomic structure.
This document outlines the key concepts to be covered in a Year 11 100 Science course on aspects of acids and bases, including atomic structure, properties of acids and bases, rates of reaction and particle theory, uses of acids and bases, and restrictions on the acids and bases included in the course. Students will study electron configuration, ionic bonding, naming ionic compounds, properties of acids and bases such as releasing hydrogen ions in water and reacting to form salts, and the rates of reactions and particle theory explanations. Assessment will include selected aspects of acids and bases such as atomic structure, properties, uses, and rates of reaction.
This document discusses the molecular structure of atoms. It begins by explaining that all substances are composed of elements, which are the basic substances, and that elements are made up of very small particles called atoms. It then describes the basic structure of an atom, including that atoms have a small, positively charged nucleus surrounded by negatively charged electrons. The number of protons determines the element and equals the number of electrons. Atoms can gain, lose, or share electrons to attain stable electron configurations, forming ions or molecules.
- Intramolecular forces are attractions between atoms within the same molecule, while intermolecular forces are attractions between different molecules.
- Covalent bonds form when atoms share electrons, and can be polar or non-polar depending on differences in electronegativity between the bonded atoms. Polar covalent bonds have partial charges while non-polar covalent bonds share electrons equally.
- The polarity of a molecule depends on differences in electronegativity between its bonded atoms, with larger differences resulting in more polar bonds and molecules. Water is an example of a polar molecule due to differences in electronegativity between oxygen and hydrogen.
1. Ionic compounds form when a metal reacts with a non-metal, resulting in positively charged metal ions and negatively charged non-metal ions that bond together in a crystalline lattice structure.
2. When ionic compounds dissolve in water or melt, the ions become free to move and conduct electricity. During electrolysis, positively charged metal ions move to the cathode and negatively charged non-metal ions move to the anode.
3. Common ionic compounds include sodium chloride, formed from sodium and chlorine ions, and copper chloride, used in electrolysis to extract copper metal from its ionic form.
This document outlines key concepts about atomic structure including:
1) The structure of atoms consists of a nucleus containing protons and neutrons surrounded by electrons. Protons are positively charged, neutrons have no charge, and electrons are negatively charged.
2) Atomic number refers to the number of protons in an atom. Mass number refers to the total number of protons and neutrons. Atomic symbols represent these numbers.
3) Electrons surround the nucleus in fixed shells. The number of electrons in each shell is limited. The arrangement of electrons is known as the electronic structure or electron configuration.
This document discusses atomic structure and electron configuration. It contains the following key points:
1. Atoms are made up of protons, neutrons, and electrons, with protons and neutrons located in the nucleus and electrons orbiting the nucleus.
2. Electrons can occupy different energy levels characterized by quantum numbers like principal quantum number and azimuthal quantum number. This determines the shape of electron orbitals.
3. The electron configuration of an atom shows the distribution of electrons among these orbitals, following the Pauli exclusion principle. Valence electrons are in the outermost shell and influence chemical bonding and properties.
4. Elements are arranged in the periodic table based on their atomic structure, including number
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D4_D5_Periodicity_Chemistry_8_pgs.docx
1. Periodicity Chemistry Worksheet
A. Periodic table
1. Which are metals? Circle your answers: C, Na, F, Cs, Ba, Ni
Which metal in the list above has the most metallic character? Explain.
2. Write the charge that each of the following atoms will have when it has a complete set of
valence electrons forming an ion.
O Na F N Ca Ar
3. What is the most common oxidation number for calcium? Explain.
Name two more elements with that oxidation number and explain your choice.
4. What element in period 3 is a metalloid? _____________________
5. When element with atomic number 118 is discovered, what family will it be in? ____________
Use the following word bank for questions 6 through 13
(Obviously, you will need to use several words more than once! )
Alkali metals Alkaline earth metals Halogens
Noble gases Transition metals Noble gases
6. The ___________________________ have a single electron in the highest energy level.
7. The ___________________________ achieve the electron configurations of noble gases by
losing two electrons.
8. The ______________________ are metals that can hold up to 10 electrons in their sublevel
shape
9. The ________________________ achieve the electron configuration of noble gases by
gaining one electron.
10. The ___________________ have full s and p orbitals in the highest occupied energy levels.
11. The ________________________ are stable and un-reactive.
12. The __________________________ are highly reactive and readily form salts with metals.
13. The _________________________ are metals that are more reactive than the transition
elements but less reactive than the alkali metals.
14. Predict the oxidation number based on the electron configuration shown.
1s2
2s2
2p6
3s2
____________ 1s2
2s2
2p6
3s1
_____________
1s2
2s2
2p6
____________ 1s2
2s2
2p5
______________
1s2
2s2
2p1
____________
2. Periodicity Chemistry Worksheet - page 2
B. Ionization Energy
1. Choose the element with the greatest first ionization energy:
Carbon or aluminum Calcium or strontium Helium or lithium
Chlorine or argon Chlorine or fluorine Sulfur or chlorine
2. Which has the larger ionization energy. sodium or potassium? Why?
3. Explain the difference in first ionization energy between lithium and beryllium.
4. The first and second ionization energies of magnesium are both relatively low, but the
third ionization energy requirement jumps to five times the previous level. Explain.
What is the most likely ion for magnesium to become when it is ionized?
5. Compare the first ionization energies for the noble gases.
6. Compare the first ionization energies for a noble gas with that of a halogen in the
same period.
7. Where would the largest jump in ionization energies be for oxygen? (with the loss of
how many electrons?)
8. How can you tell from a list of ionization energies for an element where a kernel (non
valence) electron has been removed?
3. Periodicity Chemistry Worksheet - page 3
C. Electronegativity and Electron Affinity
1. Arrange the following elements in order of increasing electronegativity.
a. gallium, aluminum, indium
b. calcium, selenium, arsenic
c. oxygen, fluorine, sulfur
d. phosphorus, oxygen, germanium
2. Will the electronegativity of barium be larger or smaller than that of strontium? Explain.
3. Compare the electronegativity of tellurium to that of antimony. Explain your reasoning.
4. The family within any period with the greatest negative electron affinity is usually the _____.
a. alkali metals b. transition metal c. halogens d. noble gases
5. Contrast ionization energy and electron affinity. In general, what can you say about these
values for metals and non-metals?
6. What is the difference between electron affinity and electronegativity?
7. If an element has a “large negative” electron affinity number where would it be located on the
periodic table?
4. Periodicity Chemistry Worksheet - page 4
D. Definitions - match
Atomic radius First ionization energy Noble gases
Decrease Increase Nonmetals
Electron affinity Ionization energy Metalloid
Electronegativity Metals Shielding effect
Noble gas configuration
1. _______ _________________ is the energy required to remove an electron from an atom.
2. The energy change associated with the addition of electron is called ________ _________.
3. The energy needed to remove the most loosely held electron from a neutral atom is called
____________________.
4. When they have a(n) ___________________, ions have a stable, filled outer electron level.
5. Along with the increased distance of the outer electrons from the nucleus, the ______
_____of the inner electrons causes ionization energy to decrease going down a column of
the periodic table.
6. A low ionization energy is characteristic of a(n) ___________________________________.
7. Ionization energies tend to _______________________ across periods of the periodic table.
8. An element with a high ionization energy is classified as a (n) ________________________.
9. The attraction an atom has for electrons is called _________________________________.
10.The distance from the nucleus to the outer most electron is known as _________________.
11.The ______________________ do not have measured electronegativites since they do not
commonly form compounds.
12.The electron arrangement with a complete outermost s and p sublevel is known
as_____________.
E. Trend Chart Draw in the trends on the periodic table:
Ionization energy electronegativity atomic radius electron affinity shielding effect
5. Periodicity Chemistry Worksheet - page 5
F. Atomic Radius
1. Circle the atom in each pair with the larger atomic radius?
Li or K Ca or Ni Ga or B O or C
Cl or Br Be or Ba Si or S Fe or Au
2. Chlorine, selenium, and bromine are located near each other on the periodic table. Which of these
elements is the smallest atom and which has the highest ionization energy?
3. Which of the following atoms is smallest: nitrogen, phosphorus, or arsenic? Which of these atoms has
the most negative electron affinity?
4. Which of the following is the largest: a potassium atom, a potassium ion with a charge of 1+ or a
rubidium atom?
5. Which of the following is the largest: a chlorine atom, a chlorine ion with a charge of 1- or a bromine
atom?
6. Which of the following is the smallest: a lithium atom, a lithium ion with a charge of 1+ or a sodium
atom?
7. Explain why within a family such as the halogens, the ionic radius increases as the atomic number
increases.
8. In terms of electron configuration and shielding, why is the atomic radius of sodium smaller than that
of potassium?
9. In terms of electron configurations and shielding, why do atoms get smaller as you move across a
period?
6. Periodicity Chemistry Worksheet - page 6
G. Concept Mastery Questions
1. The shielding effect increases with increasing atomic number within a ___.
a. period b. group c. both d. neither
2. In any ___, the number of electrons between the nucleus and the outer energy level is the same.
a. period b. group c. both d. neither
3. Within a ____, the nucleus has a stronger ability to pull on the outermost (valence) electrons in
elements of high atomic number.
a. period b. group c. both d. neither
4. In a ____, electron affinity values become more negative as atomic number increases.
a. period b. group c. both d. neither
5. The halogens are considered a ____.
a. period b. group c. both d. neither
6. Which atom has the greater nuclear charge? ______
a. Na b. Al c. P d. Ar
7. Which atom demonstrates the greatest shielding effect? _______
a. Na b. Al c. P d. Ar
8. The atoms Na, Al, P, and Ar all have the same __________
a. number of valence electrons b. size atomic radius c. number of kernel electrons
9. Which element on the periodic table has
a. lowest ionization energy __________
b. highest second ionization energy __________
c. highest electronegativity __________
d. highest ionization energy __________
e. largest atomic radius __________
10. Explain the relationship between the relative size of an cation and anion (ionic radius) to its atom
(atomic radius).
11. Explain why noble gases are inert and do not form ions.
12. Will the shielding effect be more noticeable in metals or non metals? Explain your answer.
7. Periodicity Chemistry Worksheet - page 7
13. Why do elements in the same family generally have similar properties? Choose one as an example to
support your reasoning.
14. Arrange each of the following in order of increasing ionization energy and explain your reasoning:
Calcium, iron, copper, bromine and krypton.
15. Factors affecting ionization energy include effective nuclear charge, the shielding effect, the atomic
radius and the electron arrangement in a sublevel. Use the appropriate factors to explain the overall
trend indicated by the dark line and the exceptions to it.
1) Effective Nuclear energy: Increase, Decrease or Constant
2) Shielding: Increase, Decrease or Constant
3) Atomic Radius: Increase, Decrease or Constant
4) Explain exceptions to the overall trend based
on electron configuration. _______________________
____________________________________________
16. What element am I?. (Brief periodic table location description for each clue)
Clue #1)
I have a high electron affinity, (highly negative value), and my atomic number is X.
Clue #2)
The element with atomic number X-1 has a lower ionization energy and a lower electron affinity.
Clue #3)
The element with atomic number X+1 has a higher ionization energy and basically no electron
affinity (positive value).
d) Within my group, I have the second highest ionization energy.
What element am I? ___________________________
17. What do transition metals have in common with respect to their electron configurations?
8. Periodicity Chemistry Worksheet - page 8
18. Consider the table of the first four ionization energies for an element we will call A.
Ionization Energy
(kJ/mole)
1st
2nd
3rd
4th
576 1817 2745 11580
a. In which group does A appear on the periodic table?
b. What is the most likely oxidation number for element A?
c. What is the minimum number of electrons that A must have?
d. Write the valence electron sublevel configuration for this element.
(Hint: use “n” as the energy level)
19. Can anions of two different elements have the same valence electron arrangement? If so, give
examples and discuss. If no, explain why not.
20. The first ionization energy of beryllium is 9.322 eV, the second ionization energy is 18.211 eV, and
the third ionization energy is 153.893 eV. Explain why the third ionization energy of beryllium so
much higher than the first two.