Chemical bonds are attractive forces that hold groups of atoms together. There are two main types of bonds: ionic and covalent. Ionic bonds form between metals and nonmetals via electrostatic attraction as electrons are transferred. Covalent bonds form when atoms share electrons. Lewis structures use dots and lines to represent valence electrons and show how atoms bond to achieve stable electron configurations like the octet rule.
The document summarizes key concepts about covalent bonding from a chemistry textbook chapter:
1) Covalent bonds form when two nonmetal atoms share one or more pairs of electrons to achieve a noble gas configuration, forming molecules like H2, O2, and CO2.
2) Molecular compounds formed by covalent bonds tend to have lower melting and boiling points than ionic compounds due to the weaker nature of the covalent bond.
3) Electron dot structures and Lewis diagrams are used to represent how atoms share electrons to form single, double or triple covalent bonds in molecules like H2O and NH3.
The document discusses the three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form between metals and nonmetals through the transfer of electrons from one atom to another. Covalent bonds form between nonmetals through the sharing of electrons between atoms. Metallic bonds form between metal atoms through the delocalization of valence electrons that are not assigned to any single atom. These different bonding types result in compounds and materials with distinct properties depending on the structure and interactions of the bonded atoms.
This document provides an overview of physical science concepts related to subatomic particles, the periodic table, oxidation numbers, and Lewis dot diagrams. The key points are:
1) It reviews subatomic particles, periods and groups of the periodic table, and electron configurations.
2) The objective is to predict oxidation numbers and draw Lewis dot diagrams by understanding valence electrons.
3) It defines valence electrons, oxidation numbers, and ionic bonds in chemical compounds.
4) Examples are given of writing chemical symbols, determining valence electrons, and drawing Lewis dot diagrams using the cross method.
This chapter discusses ionic bonding. It begins by explaining that noble gases have a stable electronic configuration with a full outer shell. Most other elements form ions by gaining or losing electrons to achieve a noble gas configuration. Ionic bonds form when metals transfer electrons to non-metals to form positively charged cations and negatively charged anions. Ionic compounds have a crystalline structure where the ions are arranged in a repeating pattern with strong electrostatic forces between them. This results in ionic compounds having high melting points, being insoluble in organic solvents, and capable of conducting electricity when molten or dissolved in water.
This document summarizes various chemistry concepts related to bonding:
1) Atoms bond through ionic bonding, where ions with opposite charges attract, or covalent bonding, where electrons are shared between atoms.
2) Ionic bonds form between ions, while covalent bonds form when atoms share electrons to achieve stable full outer energy levels.
3) Bonding diagrams like Lewis structures are used to represent how atoms bond by sharing or transferring electrons to achieve stable configurations.
This document provides an overview of chemical bonding and the properties of ionic and covalent compounds. It discusses the following key points:
1. Chemical bonds form due to the attraction between atoms and involve the transfer or sharing of valence electrons. Ionic bonds form through electron transfer between metals and nonmetals, while covalent bonds involve electron sharing.
2. Lewis symbols represent atoms and their valence electrons and are used to predict bonding patterns. Electronegativity determines bond polarity.
3. Ionic compounds have high melting and boiling points due to strong electrostatic attractions in the crystal lattice. Covalent compounds can be solids, liquids or gases.
1) Chemical bonds form when atoms overlap their orbitals to achieve stable noble gas configurations. This increases stability as atoms form ionic or covalent bonds.
2) Metallic bonding occurs via a "sea of electrons" model where mobile electrons are shared between rigid positive ions. This explains properties like conductivity and malleability.
3) There are various types of bonds including ionic formed between metals and nonmetals, and covalent including polar, nonpolar, and coordinate bonds formed by electron sharing or donation.
The document discusses Lewis structures and the rules for drawing them. It explains that Lewis structures show how atoms bond via shared electron pairs to achieve stable noble gas configurations. It provides a 4-step process for drawing Lewis structures, covering counting electrons, identifying the central atom, adding lone pairs to complete octets, and checking that all electrons are accounted for. Exceptions to the octet rule and drawing structures for ions are also covered.
The document summarizes key concepts about covalent bonding from a chemistry textbook chapter:
1) Covalent bonds form when two nonmetal atoms share one or more pairs of electrons to achieve a noble gas configuration, forming molecules like H2, O2, and CO2.
2) Molecular compounds formed by covalent bonds tend to have lower melting and boiling points than ionic compounds due to the weaker nature of the covalent bond.
3) Electron dot structures and Lewis diagrams are used to represent how atoms share electrons to form single, double or triple covalent bonds in molecules like H2O and NH3.
The document discusses the three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form between metals and nonmetals through the transfer of electrons from one atom to another. Covalent bonds form between nonmetals through the sharing of electrons between atoms. Metallic bonds form between metal atoms through the delocalization of valence electrons that are not assigned to any single atom. These different bonding types result in compounds and materials with distinct properties depending on the structure and interactions of the bonded atoms.
This document provides an overview of physical science concepts related to subatomic particles, the periodic table, oxidation numbers, and Lewis dot diagrams. The key points are:
1) It reviews subatomic particles, periods and groups of the periodic table, and electron configurations.
2) The objective is to predict oxidation numbers and draw Lewis dot diagrams by understanding valence electrons.
3) It defines valence electrons, oxidation numbers, and ionic bonds in chemical compounds.
4) Examples are given of writing chemical symbols, determining valence electrons, and drawing Lewis dot diagrams using the cross method.
This chapter discusses ionic bonding. It begins by explaining that noble gases have a stable electronic configuration with a full outer shell. Most other elements form ions by gaining or losing electrons to achieve a noble gas configuration. Ionic bonds form when metals transfer electrons to non-metals to form positively charged cations and negatively charged anions. Ionic compounds have a crystalline structure where the ions are arranged in a repeating pattern with strong electrostatic forces between them. This results in ionic compounds having high melting points, being insoluble in organic solvents, and capable of conducting electricity when molten or dissolved in water.
This document summarizes various chemistry concepts related to bonding:
1) Atoms bond through ionic bonding, where ions with opposite charges attract, or covalent bonding, where electrons are shared between atoms.
2) Ionic bonds form between ions, while covalent bonds form when atoms share electrons to achieve stable full outer energy levels.
3) Bonding diagrams like Lewis structures are used to represent how atoms bond by sharing or transferring electrons to achieve stable configurations.
This document provides an overview of chemical bonding and the properties of ionic and covalent compounds. It discusses the following key points:
1. Chemical bonds form due to the attraction between atoms and involve the transfer or sharing of valence electrons. Ionic bonds form through electron transfer between metals and nonmetals, while covalent bonds involve electron sharing.
2. Lewis symbols represent atoms and their valence electrons and are used to predict bonding patterns. Electronegativity determines bond polarity.
3. Ionic compounds have high melting and boiling points due to strong electrostatic attractions in the crystal lattice. Covalent compounds can be solids, liquids or gases.
1) Chemical bonds form when atoms overlap their orbitals to achieve stable noble gas configurations. This increases stability as atoms form ionic or covalent bonds.
2) Metallic bonding occurs via a "sea of electrons" model where mobile electrons are shared between rigid positive ions. This explains properties like conductivity and malleability.
3) There are various types of bonds including ionic formed between metals and nonmetals, and covalent including polar, nonpolar, and coordinate bonds formed by electron sharing or donation.
The document discusses Lewis structures and the rules for drawing them. It explains that Lewis structures show how atoms bond via shared electron pairs to achieve stable noble gas configurations. It provides a 4-step process for drawing Lewis structures, covering counting electrons, identifying the central atom, adding lone pairs to complete octets, and checking that all electrons are accounted for. Exceptions to the octet rule and drawing structures for ions are also covered.
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to achieve stable octets. Atoms form covalent bonds by sharing electrons, depicted using Lewis dot structures. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals.
1) Chemical bonds can be either ionic or covalent. Ionic bonds form when electrons are transferred between metals and non-metals to form ions. Covalent bonds form when electrons are shared between non-metals.
2) Sodium chloride forms when sodium donates an electron to chlorine to form ions that are attracted in an ionic bond. Hydrogen molecule forms when hydrogen atoms share an electron pair in a single covalent bond.
3) Ionic compounds have high melting points, conduct electricity when molten or dissolved, and dissolve in water but not organic solvents. Covalent compounds have lower melting points, do not conduct electricity, and dissolve in organic solvents but not water.
This document provides an overview of ionic and covalent bonding. It discusses the formation of ions through the loss or gain of valence electrons to achieve stable electron configurations. Ionic compounds are formed between metallic and nonmetallic elements and are held together by ionic bonds between cations and anions. Molecular compounds are formed by the sharing of valence electrons between nonmetallic elements to form covalent bonds. Polar and nonpolar covalent bonds are discussed based on differences in electronegativity between bonded atoms. Hydrogen bonds that occur between polar molecules like water are also summarized. Key terms related to ionic bonding, covalent bonding, and molecular structure are defined.
The document discusses the structure and bonding of atoms and molecules. It begins by describing the components of an atom, including protons, neutrons, and electrons. It then discusses the periodic table and how elements in the same row or column have similar properties. The document goes on to describe atomic orbitals like s and p orbitals. It also discusses how elements bond, including ionic and covalent bonding. Additional topics covered include Lewis structures, resonance structures, molecular geometry, and organic naming conventions.
This document provides an overview of key concepts in chemistry including:
1) The structure of atoms including protons, neutrons, and electrons. It also discusses isotopes and electron configuration.
2) The periodic table is introduced including periodic trends in properties and how elements are arranged in groups and periods. Metals, nonmetals, and chemical properties are also covered.
3) Bonding including ionic bonding between metals and nonmetals and covalent bonding between nonmetals is explained through examples like sodium chloride and water. Dot and cross diagrams are used to represent covalent bonds.
4) Compounds and chemical equations are discussed including balancing equations and calculating relative formula mass. Giant ionic structures
Ionic bonding occurs between metal and non-metal atoms when they form ions. Metals form positive ions by losing electrons, filling their outer electron shells. Non-metals form negative ions by gaining electrons. The oppositely charged ions are attracted in an ionic compound via electrostatic forces. Sodium chloride is an example where sodium atoms lose electrons to become Na+ ions and chloride atoms gain electrons to become Cl- ions. The ions are arranged in a crystal lattice structure held together by ionic bonds.
The document discusses ionic bonding between metals and non-metals. Ionic bonding occurs when metals give up electrons to form positive ions and non-metals gain electrons to form negative ions. The positive and negative ions are then attracted to each other, forming an ionic compound. Metals are usually found on the left side of the periodic table and easily give up valence electrons. Non-metals are usually on the right side and readily gain electrons to achieve a full valence shell. When ions form, they arrange in a crystalline lattice structure with positive and negative ions alternating. Ionic compounds have properties like being crystalline solids, having high melting points, and being able to conduct electricity when melted or
An electron is a negatively charged particle that orbits the nucleus of an atom. A proton is a positively charged particle found within the nucleus. A neutron is a particle within the nucleus that has no charge. Atoms are made up of electrons orbiting a nucleus containing protons and neutrons. Elements are substances made of only one type of atom that cannot be broken down further by chemical or physical changes. Atoms of the same element have the same number of protons but can differ in the number of neutrons, forming isotopes of that element.
The document discusses atomic structure and mass spectrometry. It defines key terms like mass number, atomic number, and isotope. It explains the process of mass spectrometry, including ionization, acceleration, deflection, and detection of ions. Graphs of ionization energies are analyzed to determine electronic configurations and periodic trends. Successive ionization energies are explained by electron shielding effects. Radioactive decay and half-life are also defined.
This document summarizes key concepts from a chemistry textbook chapter on the periodic table. It discusses how elements are organized in the periodic table based on their atomic structure and properties. Early sections describe the historical development of periodic tables and how elements are classified into groups based on electron configuration. Later sections summarize periodic trends in atomic size, ionization energy, and ion size based on an element's position in the periodic table and how its electron configuration is filled.
This document summarizes key concepts about atoms and molecules. It defines an atom as the smallest unit that retains an element's characteristics. Atoms have a nucleus containing protons and neutrons surrounded by an electron cloud. Molecules are defined as stable groups of two or more atoms bonded together, either through ionic bonds formed by electron transfer or covalent bonds formed by electron sharing. Examples of hydrogen, lithium, and argon atoms are provided to illustrate their atomic structure. Isotopes are also introduced as atoms of the same element with different numbers of neutrons.
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Chemistry is the study of matter and its changes. An atom is the smallest particle of an element that retains the properties of that element. Atoms combine to form compounds with fixed ratios. A chemical change alters the composition of a substance, while a physical change does not. The atomic theory states that elements are composed of atoms and compounds are composed of two or more different elements chemically bonded. The structure of the atom consists of a small, dense nucleus surrounded by electrons. Chemical formulas represent the elements and their ratios in compounds and molecules. Ions are formed when atoms gain or lose electrons. Naming and formulas help identify substances. Acids donate hydrogen ions in water and bases donate hydroxide ions.
This document provides an overview of chemical bonding concepts including:
- The octet rule which states that main group elements form ions to achieve 8 valence electrons.
- Ionic and covalent bonds are formed through the transfer or sharing of electrons respectively.
- Lewis structures are used to represent electron pairing in molecules and predict molecular geometry based on electron pair repulsion.
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The document outlines key learning outcomes and concepts about atomic structure, including describing the structure of atoms with atomic numbers 1 to 20, defining terms like atomic number and mass number, explaining electron configuration and outer electrons, and distinguishing between isotopes, ions, and molecules of elements and compounds. It also provides illustrations of atomic structure and examples of applying atomic structure concepts.
Covalent bonds form when two nonmetal atoms share one or more pairs of valence electrons to achieve a noble gas configuration. This sharing of electrons results in the formation of molecules with lower potential energy than the individual atoms. Covalent bonds can be either polar or nonpolar depending on whether the electrons are shared equally or unequally between the bonded atoms. Multiple bonds are also possible where two or more pairs of electrons are shared, such as double and triple bonds.
Valence electrons are the outermost shell electrons of an atom that are involved in bonding. Elements in the same group on the periodic table have the same number of valence electrons because they exhibit similar chemical properties based on their valence electron configuration. Atoms seek to attain a full outer shell of 8 electrons to achieve stability through gaining, losing or sharing valence electrons in chemical bonds.
Ionic bonds form when oppositely charged ions attract each other, forming ionic compounds. Cations form when atoms lose electrons to achieve a stable electron configuration, while anions form when atoms gain electrons. Ionic compounds consist of a crystal lattice structure where cations are surrounded by anions. They have properties like high melting points and boiling points since energy is required to overcome the strong electrostatic attractions between ions.
There are three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form when a metal transfers electrons to a nonmetal, creating positively charged ions and negatively charged ions. Covalent bonds form when atoms share electrons as either single, double or triple bonds. Metallic bonds form a "sea of electrons" that are shared between positive metal ions throughout a crystalline structure.
This document discusses different types of chemical bonds including ionic, covalent, and metallic bonds. It describes the concepts of electronegativity, resonance structures, and exceptions to the octet rule. Bond strength is quantified by bond enthalpy, which is the energy required to break a bond. Stronger bonds have higher bond enthalpies and shorter bond lengths.
The document provides information about chemical bonding and different types of bonds. It begins by defining a chemical bond as the forces that hold groups of atoms together, and explains that bonds form when the energy of bonded atoms is lower than separated atoms. It then describes the main types of bonds:
- Ionic bonds result from the transfer of electrons between metals and nonmetals.
- Covalent bonds result from the sharing of electrons between atoms.
- Polar covalent bonds occur when electrons are unequally shared, resulting in partial charges.
The document discusses electronegativity and how it relates to bond polarity. It also introduces dipole moments and how bond polarity affects molecular properties like solubility. Finally, it explains
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to achieve stable octets. Atoms form covalent bonds by sharing electrons, depicted using Lewis dot structures. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals.
1) Chemical bonds can be either ionic or covalent. Ionic bonds form when electrons are transferred between metals and non-metals to form ions. Covalent bonds form when electrons are shared between non-metals.
2) Sodium chloride forms when sodium donates an electron to chlorine to form ions that are attracted in an ionic bond. Hydrogen molecule forms when hydrogen atoms share an electron pair in a single covalent bond.
3) Ionic compounds have high melting points, conduct electricity when molten or dissolved, and dissolve in water but not organic solvents. Covalent compounds have lower melting points, do not conduct electricity, and dissolve in organic solvents but not water.
This document provides an overview of ionic and covalent bonding. It discusses the formation of ions through the loss or gain of valence electrons to achieve stable electron configurations. Ionic compounds are formed between metallic and nonmetallic elements and are held together by ionic bonds between cations and anions. Molecular compounds are formed by the sharing of valence electrons between nonmetallic elements to form covalent bonds. Polar and nonpolar covalent bonds are discussed based on differences in electronegativity between bonded atoms. Hydrogen bonds that occur between polar molecules like water are also summarized. Key terms related to ionic bonding, covalent bonding, and molecular structure are defined.
The document discusses the structure and bonding of atoms and molecules. It begins by describing the components of an atom, including protons, neutrons, and electrons. It then discusses the periodic table and how elements in the same row or column have similar properties. The document goes on to describe atomic orbitals like s and p orbitals. It also discusses how elements bond, including ionic and covalent bonding. Additional topics covered include Lewis structures, resonance structures, molecular geometry, and organic naming conventions.
This document provides an overview of key concepts in chemistry including:
1) The structure of atoms including protons, neutrons, and electrons. It also discusses isotopes and electron configuration.
2) The periodic table is introduced including periodic trends in properties and how elements are arranged in groups and periods. Metals, nonmetals, and chemical properties are also covered.
3) Bonding including ionic bonding between metals and nonmetals and covalent bonding between nonmetals is explained through examples like sodium chloride and water. Dot and cross diagrams are used to represent covalent bonds.
4) Compounds and chemical equations are discussed including balancing equations and calculating relative formula mass. Giant ionic structures
Ionic bonding occurs between metal and non-metal atoms when they form ions. Metals form positive ions by losing electrons, filling their outer electron shells. Non-metals form negative ions by gaining electrons. The oppositely charged ions are attracted in an ionic compound via electrostatic forces. Sodium chloride is an example where sodium atoms lose electrons to become Na+ ions and chloride atoms gain electrons to become Cl- ions. The ions are arranged in a crystal lattice structure held together by ionic bonds.
The document discusses ionic bonding between metals and non-metals. Ionic bonding occurs when metals give up electrons to form positive ions and non-metals gain electrons to form negative ions. The positive and negative ions are then attracted to each other, forming an ionic compound. Metals are usually found on the left side of the periodic table and easily give up valence electrons. Non-metals are usually on the right side and readily gain electrons to achieve a full valence shell. When ions form, they arrange in a crystalline lattice structure with positive and negative ions alternating. Ionic compounds have properties like being crystalline solids, having high melting points, and being able to conduct electricity when melted or
An electron is a negatively charged particle that orbits the nucleus of an atom. A proton is a positively charged particle found within the nucleus. A neutron is a particle within the nucleus that has no charge. Atoms are made up of electrons orbiting a nucleus containing protons and neutrons. Elements are substances made of only one type of atom that cannot be broken down further by chemical or physical changes. Atoms of the same element have the same number of protons but can differ in the number of neutrons, forming isotopes of that element.
The document discusses atomic structure and mass spectrometry. It defines key terms like mass number, atomic number, and isotope. It explains the process of mass spectrometry, including ionization, acceleration, deflection, and detection of ions. Graphs of ionization energies are analyzed to determine electronic configurations and periodic trends. Successive ionization energies are explained by electron shielding effects. Radioactive decay and half-life are also defined.
This document summarizes key concepts from a chemistry textbook chapter on the periodic table. It discusses how elements are organized in the periodic table based on their atomic structure and properties. Early sections describe the historical development of periodic tables and how elements are classified into groups based on electron configuration. Later sections summarize periodic trends in atomic size, ionization energy, and ion size based on an element's position in the periodic table and how its electron configuration is filled.
This document summarizes key concepts about atoms and molecules. It defines an atom as the smallest unit that retains an element's characteristics. Atoms have a nucleus containing protons and neutrons surrounded by an electron cloud. Molecules are defined as stable groups of two or more atoms bonded together, either through ionic bonds formed by electron transfer or covalent bonds formed by electron sharing. Examples of hydrogen, lithium, and argon atoms are provided to illustrate their atomic structure. Isotopes are also introduced as atoms of the same element with different numbers of neutrons.
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Chemistry is the study of matter and its changes. An atom is the smallest particle of an element that retains the properties of that element. Atoms combine to form compounds with fixed ratios. A chemical change alters the composition of a substance, while a physical change does not. The atomic theory states that elements are composed of atoms and compounds are composed of two or more different elements chemically bonded. The structure of the atom consists of a small, dense nucleus surrounded by electrons. Chemical formulas represent the elements and their ratios in compounds and molecules. Ions are formed when atoms gain or lose electrons. Naming and formulas help identify substances. Acids donate hydrogen ions in water and bases donate hydroxide ions.
This document provides an overview of chemical bonding concepts including:
- The octet rule which states that main group elements form ions to achieve 8 valence electrons.
- Ionic and covalent bonds are formed through the transfer or sharing of electrons respectively.
- Lewis structures are used to represent electron pairing in molecules and predict molecular geometry based on electron pair repulsion.
Sue welcomes viewers to her new show. The show stars Sue and will feature her talents and skills. Sue is excited to entertain audiences and show off her abilities on television. Viewers are sure to enjoy Sue's unique personality and performance.
The document outlines key learning outcomes and concepts about atomic structure, including describing the structure of atoms with atomic numbers 1 to 20, defining terms like atomic number and mass number, explaining electron configuration and outer electrons, and distinguishing between isotopes, ions, and molecules of elements and compounds. It also provides illustrations of atomic structure and examples of applying atomic structure concepts.
Covalent bonds form when two nonmetal atoms share one or more pairs of valence electrons to achieve a noble gas configuration. This sharing of electrons results in the formation of molecules with lower potential energy than the individual atoms. Covalent bonds can be either polar or nonpolar depending on whether the electrons are shared equally or unequally between the bonded atoms. Multiple bonds are also possible where two or more pairs of electrons are shared, such as double and triple bonds.
Valence electrons are the outermost shell electrons of an atom that are involved in bonding. Elements in the same group on the periodic table have the same number of valence electrons because they exhibit similar chemical properties based on their valence electron configuration. Atoms seek to attain a full outer shell of 8 electrons to achieve stability through gaining, losing or sharing valence electrons in chemical bonds.
Ionic bonds form when oppositely charged ions attract each other, forming ionic compounds. Cations form when atoms lose electrons to achieve a stable electron configuration, while anions form when atoms gain electrons. Ionic compounds consist of a crystal lattice structure where cations are surrounded by anions. They have properties like high melting points and boiling points since energy is required to overcome the strong electrostatic attractions between ions.
There are three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form when a metal transfers electrons to a nonmetal, creating positively charged ions and negatively charged ions. Covalent bonds form when atoms share electrons as either single, double or triple bonds. Metallic bonds form a "sea of electrons" that are shared between positive metal ions throughout a crystalline structure.
This document discusses different types of chemical bonds including ionic, covalent, and metallic bonds. It describes the concepts of electronegativity, resonance structures, and exceptions to the octet rule. Bond strength is quantified by bond enthalpy, which is the energy required to break a bond. Stronger bonds have higher bond enthalpies and shorter bond lengths.
The document provides information about chemical bonding and different types of bonds. It begins by defining a chemical bond as the forces that hold groups of atoms together, and explains that bonds form when the energy of bonded atoms is lower than separated atoms. It then describes the main types of bonds:
- Ionic bonds result from the transfer of electrons between metals and nonmetals.
- Covalent bonds result from the sharing of electrons between atoms.
- Polar covalent bonds occur when electrons are unequally shared, resulting in partial charges.
The document discusses electronegativity and how it relates to bond polarity. It also introduces dipole moments and how bond polarity affects molecular properties like solubility. Finally, it explains
This chapter discusses chemical bonds and mixtures. It introduces electron-dot structures to show valence electrons and how they are involved in bonding. Ionic bonds form when ions with opposite charges are attracted to each other. Covalent bonds form when atoms share electrons. Polar covalent bonds result when electrons are shared unevenly. Molecular polarity arises if the polar bonds in a molecule do not cancel out. Most materials are mixtures that can be separated into pure substances. Solutions are homogeneous mixtures where one substance dissolves evenly throughout another. Concentration, molarity, and solubility are measures used to describe solutions.
The document discusses chemical bonding and molecular structures. It explains that chemical bonding occurs through ionic bonding via the transfer of electrons between atoms, or covalent bonding via the sharing of electron pairs between atoms. It also describes molecular geometry models including VSEPR theory, which predicts the three-dimensional arrangements of atoms in molecules based on electron pair repulsion. Common molecular shapes such as linear, trigonal planar, tetrahedral and octahedral are defined.
1) Atoms form chemical bonds in order to attain a stable electron configuration with 8 valence electrons, known as an octet.
2) There are three main types of bonds: ionic bonds form when atoms transfer electrons to become ions, covalent bonds form when atoms share electrons, and metallic bonds involve delocalized electrons distributed among positively charged metal ions.
3) Whether a bond is ionic or covalent depends on the electronegativity of the atoms involved - ionic bonds form between metals and nonmetals, while covalent bonds form between two nonmetals.
This document discusses three main types of chemical bonding: ionic bonds, covalent bonds, and metallic bonds. Ionic bonds form between metals and nonmetals through the transfer of electrons from one atom to another, producing charged ions. Covalent bonds form between nonmetals of similar electronegativity through the sharing of electron pairs. Metallic bonds form between metallic elements through a shared "electron cloud" that holds the atoms together strongly.
Covalent bonds form between nonmetal atoms by sharing valence electrons. Atoms share electrons to attain stable electron configurations like noble gases. Lewis structures show how valence electrons are arranged between bonded atoms. To draw Lewis structures, count the total valence electrons and distribute them to form single or double bonds between atoms until each atom has an octet of electrons. Examples of molecules held by covalent bonds are hydrogen, oxygen, and chlorine.
Chemical bonds form when atoms share or transfer electrons. There are several main types of bonds:
- Ionic bonds form when metals transfer electrons to nonmetals to form positive and negative ions that are attracted to each other. Ionic compounds are crystalline and dissolve in water.
- Covalent bonds form when atoms share two or more valence electrons to achieve stability. Covalent bond strength depends on the number of electron pairs shared. Covalent compounds exist as discrete molecules.
- Metallic bonds result from the attraction between positively charged metal ions and delocalized electrons in the "sea of electrons" in the solid metal. Metallic bonding explains the properties of metals like conductivity.
Atoms form bonds by gaining, losing, or sharing electrons to attain a stable electron configuration with 8 valence electrons, known as an octet. Ionic bonds form when atoms transfer electrons to become ions with opposite charges that are attracted to each other. Covalent bonds form when atoms share electrons to attain a full outer shell. Common diatomic molecules that form through covalent bonding include hydrogen (H2), oxygen (O2), nitrogen (N2), fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2). Ionic bonds typically form between metals and nonmetals, while covalent bonds form between two nonmetals.
This document presents information about chemical bonding from the chemistry project of students Akarshik Banerjee, Pratyush Dey, and Sayantan Biswas. It discusses various types of chemical bonds including covalent bonds, which form when atoms share electron pairs, and ionic bonds, which form through complete electron transfer. It also describes concepts like the octet rule, Lewis dot structures, formal charge, resonance structures, and molecular geometry based on valence shell electron pair repulsion theory. Hybridization and molecular orbitals are explained as well. In summary, the document provides an overview of key concepts in chemical bonding from the perspective of high school chemistry students.
The document discusses atomic structure and bonding. It covers subatomic particles like protons, neutrons and electrons. It defines atomic number and atomic mass. Electrons orbit the nucleus in energy levels. Ionic bonding occurs when atoms gain or lose electrons to achieve full outer shells, forming ions. Covalent bonding occurs when atoms share electrons. Ionic bonds are usually strong solids while covalent bonds are weaker liquids or gases.
This document discusses the four main types of chemical bonds: ionic bonds, covalent bonds, hydrogen bonds, and metallic bonds. Ionic bonds involve the transfer of electrons between atoms. Covalent bonds involve the sharing of electrons between two atoms. Hydrogen bonds are electrostatic attractions between hydrogen atoms covalently bonded to electronegative atoms and another electronegative atom. Metallic bonds are electrostatic attractions between positively charged metal ions and delocalized electrons in metals. Examples of each type of bond are provided.
Chemical bonds form when atoms attract each other and bind together. There are three main types of bonds: ionic bonds form when a metal transfers electrons to a non-metal, metallic bonds involve delocalized electrons that move freely between metal atoms, and covalent bonds occur when two non-metals share pairs of electrons. Ionic bonds are strong but brittle, metallic bonds allow metals to conduct heat and electricity, and covalent bonds can be single, double or triple depending on how many electron pairs are shared.
Chemical bonds result from the attraction between positively charged nuclei and negatively charged electrons. There are three main types of bonding: ionic, metallic, and covalent. Ionic bonding involves the transfer of electrons between atoms to form ions with opposite charges that are attracted to each other. Covalent bonding involves the sharing of electrons between nonmetal atoms. Molecules form through covalent bonds and have molecular formulas that describe the number and type of atoms present. The shape and polarity of molecules can be predicted from their Lewis structures.
1. The document discusses different types of chemical formulas including molecular, empirical, and structural formulas. It provides examples of each type like H2O for water and C6H12 for hexene.
2. It also discusses ionic and covalent bonding. Ionic bonding involves the complete transfer of electrons from one atom to another, like from sodium to chlorine in NaCl. Covalent bonding involves the sharing of electron pairs between atoms.
3. The document describes electronegativity and how it relates to the polarity of covalent bonds. Polar covalent bonds form between atoms with an electronegativity difference of 0.5-1.6, while ionic bonds form between atoms with a difference above
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This document discusses covalent bonding. Covalent bonding occurs when two nonmetal atoms share pairs of electrons to acquire a stable octet of electrons. Lewis dot structures can be used to represent covalent bonds by showing how valence electrons from different atoms are shared. The shapes of covalent molecules are determined by VSEPR theory, which states that molecules will adopt a geometry that minimizes electron pair repulsion. Common molecular shapes include linear, trigonal planar, tetrahedral, bent, and trigonal pyramidal.
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to achieve stable octets. Atoms form covalent bonds by sharing electrons, depicted using Lewis dot structures. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals.
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to acquire stable electron configurations. Atoms form covalent bonds by sharing electrons, with the shared electrons shown as dots in Lewis dot structures. The three major types of bonding are ionic bonding, metallic bonding and covalent bonding. Covalent bonding forms molecules by electron sharing between nonmetal atoms.
This document provides an overview of basic chemistry concepts. It defines elements and atoms, and explains that all matter is composed of elements. Atoms are made up of subatomic particles including protons, neutrons, and electrons. The periodic table arranges elements and displays their atomic numbers and masses. Molecules are formed via chemical bonds between elements, including ionic bonds and covalent bonds. Water is a polar molecule that has important properties like being a good solvent and existing in solid, liquid, and gas forms. Acids and bases are defined based on their ability to donate or accept hydrogen ions in water. The pH scale measures hydrogen ion concentration to indicate acidity or alkalinity. Buffers help maintain pH within a normal
1. The document discusses covalent bonding and molecular compounds. It defines covalent bonds as the sharing of electrons between nonmetal atoms.
2. Molecular compounds are formed from covalent bonds between atoms. They have lower melting and boiling points than ionic compounds.
3. Molecular formulas show the number and type of atoms in a molecule, but not their arrangement. Water's molecular formula is H2O.
Covalent bonding occurs when two nonmetal atoms share pairs of electrons to acquire stable electron configurations. Atoms form covalent bonds by sharing electrons, drawing Lewis dot structures to show these electron pairs. There are three main types of bonding: ionic, metallic, and covalent. Covalent bonding is the focus of this chapter and forms molecules by electron sharing between nonmetals. Molecular compounds are named using prefixes to indicate the number of each element present. Bond polarity is determined using electronegativity values and can classify bonds and molecules as ionic, polar covalent, or nonpolar covalent.
This document provides information about molecular and ionic compounds, including:
- Molecular compounds are formed by covalent bonds between nonmetal atoms, while ionic compounds involve metal and nonmetal atoms bonded by ionic bonds.
- Molecular formulas show the actual number and type of atoms in a molecule, while ionic formulas use the lowest whole number ratio.
- Covalent bonds are represented by electron dot structures that show how atoms share electrons to achieve stable configurations. Multiple and coordinate covalent bonds are also discussed.
- Polarity arises in polar covalent bonds due to unequal electron sharing. Polar molecules have dipole moments while intermolecular forces include hydrogen bonding, dipole-dipole interactions, and
This document provides an overview of the structure of matter and different types of bonds between atoms. It defines key terms like elements, compounds, and chemical bonds. It describes the three main types of bonds - covalent, ionic, and metallic - and explains how they form. Covalent bonds form when atoms share electrons. Ionic bonds form when electrons are transferred between atoms to form ions. Metallic bonds form due to attraction between positively charged metal ions and delocalized electrons. The document also discusses naming conventions for compounds and polyatomic ions.
The document discusses covalent bonding and Lewis dot structures. It provides examples of how atoms share electrons to form covalent bonds in order to achieve stable octet configurations. Diatomic molecules such as H2, O2, N2, F2, and many biological molecules form covalent bonds in this way. Lewis dot structures are used to represent how valence electrons are arranged among atoms in molecules. Resonance structures can occur when more than one valid Lewis structure can be drawn for a molecule.
Electrons in atoms are arranged in discrete shells around the nucleus. The first shell holds 2 electrons and subsequent shells hold up to 8 electrons each following the octet rule. Ions form when atoms gain or lose electrons to fill or empty their outer shell, becoming negatively or positively charged. Ionic bonds occur when oppositely charged ions attract. Covalent bonds occur when atoms share electrons in molecular orbitals. Nonpolar covalent bonds share electrons equally while polar covalent bonds have an unequal electron distribution making one end partially negative and the other partially positive. Hydrogen bonds form between the partially positive hydrogen of one polar molecule and the partially negative end of another.
1. Covalent bonds form when two atoms share one or more pairs of valence electrons in order to achieve a stable octet of electrons.
2. Molecules are formed when atoms are bonded together by covalent bonds, and molecular compounds are composed of molecules.
3. Molecular compounds tend to have lower melting and boiling points than ionic compounds and many are gases or liquids at room temperature.
This document discusses the structure of matter and different types of chemical bonds. It begins by defining important terms like elements, atoms, and compounds. It then explains the three main types of atomic bonds: covalent, ionic, and metallic. Covalent bonds form when atoms share electrons, ionic bonds form when electrons are transferred between atoms to form ions, and metallic bonds result from the attraction between positively charged metal ions and delocalized electrons. The document provides examples of each bond type and discusses naming conventions for compounds. It also covers polyatomic ions and transition metals.
This document provides an overview of chemical bonding concepts including ionic bonds, covalent bonds, electronegativity, and molecular shapes. Key points covered include: 1) Ionic bonds form between cations and anions via electrostatic attraction while covalent bonds form through the sharing of electron pairs. 2) Electronegativity determines the polarity of covalent bonds, with more electronegative atoms attracting bonding electrons. 3) VSEPR theory predicts molecular geometry based on electron pair-atom repulsion.
chap8lect_2015, perteneciente a fiisca del estado solido.pptJorgespw
The document summarizes key concepts in chemical bonding, including the three main types of bonds (ionic, covalent, and metallic), ionic bonding between metals and nonmetals, energetics of ionic bonding involving ionization energy, electron affinity, and lattice energy, properties of covalent bonding including polar covalent bonds and electronegativity, Lewis structures for representing covalent bonding including exceptions to the octet rule, and resonance structures.
This document discusses different types of chemical bonds including ionic, covalent, and metallic bonds. It describes the formation of ionic bonds between metals and nonmetals and how ionization energy, electron affinity, and lattice energy contribute to the energetics of ionic bonding. Covalent bonding is explained as the sharing of electrons between nonmetals. Factors that determine bond polarity like electronegativity are also covered. The document provides details on writing Lewis structures, accounting for valence electrons and formal charges. Exceptions to the octet rule for molecules with odd numbers of electrons, incomplete octets, and expanded octets are explained.
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This document discusses chemical bonding and Lewis dot structures. It begins by listing the objectives of describing chemical bonding, representing atoms with Lewis diagrams, and applying rules to deduce bond formation. It then defines chemical bonds as electrostatic attractions between atoms that share electrons. It discusses valence and core electrons and uses oxygen as an example. It introduces Lewis structures and represents oxygen and sulfur. It explains how covalent bonds form between hydrogen atoms and represents this with an energy diagram. It provides examples of Lewis structures for various elements and molecules such as HCl, H2O, O2, and HCN. It also discusses lone pairs and coordinate covalent bonds. Finally, it provides examples and questions to solidify the concepts.
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Intro chem ch 12 chemical bonding sp08
1. 1
Introduction to Chemistry
Chapter 12
Chemical Bonds
Chemical Bonds are the attractive forces that hold groups of two or more atoms together and make them function as a unit.
Chemical bonding occurs with the valence electrons of an atom.
Two types of chemical bonds:
Ionic: Ionic bonding occurs when an atom that loses electrons relatively easily reacts with an atom that gains
electrons relatively easily.
• metals with nonmetals
• salts
• crystals
• held together by electrostatic attraction
E.g. How salts are formed:
Na0 (s) + Cl20
2 Na+ + Cl– + Cl– → NaCl + NaCl
or
2 Na+ + 2Cl– 2 NaCl
Ionic compounds: (formula units)
• are solids with high melting points
• they exist as crystal structures
• they consist of formula units
• most are soluble in water
• molten compounds conduct electricity well because they contain mobile charged ions.
• aqueous solutions (dissolved in water) conduct electricity well because they contain mobile charged ions.
• they are mostly formed between two elements with large differences (∆EN>2) in electronegativity, usually a metal
and a nonmetal.
Covalent bonding: This bonding occurs when electrons are shared by atoms. The electrons are attracted to the
nucleus of the opposing atoms in the bond.
E.g. H2 HCl SiO2
H• •H H··H H—H H2
•• •• ..
H• • Br ׃ H ׃Br ׃ H—Br : HBr
•• •• ..
Covalent compounds: (molecules)
• they are solids, liquids, and gases with low melting points
• many are not soluble in water
• liquid and molten compounds do not conduct electricity
• aqueous solutions are usually poor conductors of electricity because most do not contain charged particles (ions).
• they are often formed between two elements with similar electronegativities, usually nonmetals.
To understand chemical bonding, we must first understand the forces between the bonds.
• Electronegativity: is the ability of an atom in a molecule to attract shared electrons to itself.
2. 2
2
0 4
∆ EN
0 = covalent
> 0 = polar covalent
> 2 = ionic
∆EN = H—H 2.1-2.1 = 0 covalent
∆EN = S—H 2.5-2.1 = 0.4 polar covalent
∆EN = NaCl 0.9-3.0 = 2.1 ionic
Lewis Dots:
• Lewis dots represent the valence electrons of an atom.
• The number of valence electrons for Group A elements are the same as the group number.
3. 3
Lewis Structures: The Octet Rule
The theory about why noble gases are basically unreactive and stable was built around the observations:
• Noble gases have completely filled s and p shells
• There are 8 valence electrons for each of the noble gases
• The filled shells of the valence electrons account for the stability of the noble gases
The octet rule states that atoms will share, gain, or lose electrons in order to produce a noble gas electron configuration.
He → 1s2 Li+ → 1s2 Be2+ → 1s2
Ne → 1s22s22p6 Na+ → 1s22s22p6 Mg2+→ 1s22s22p6 Al3+→1s22s22p6
F– → 1s22s22p6 O2– → 1s22s22p6 N3– → 1s22s22p6
Isoelectronic – elements and ions that contain the same number of electrons.
Ionic Bonding
General Formula Table
4. 4
Ionic compounds: Lewis structures
..
NaCl → Na+ + [ :Cl: ]–
¨
..
MgBr2 Mg2+ + 2 [ :Br:]–
¨
..
K2S 2 K+ + [:S: ]2–
¨
..
Al2O3 2 Al3+ + 3 [:O: ]2–
¨
Lewis Structure for Covalent Compounds:
• Share electrons
• Sometimes the octet rule applies
• Carbon, nitrogen, oxygen, and fluorine always obey the octet rule
H : H The dots between the hydrogen atoms represent shared electrons.
H—H The straight line between the atoms represents two electrons shared between the two hydrogen atoms
Rules to determine the number of bonds needed in covalent bonding:
5. 5
Binary compounds:
1. Determine the number of valence electrons for each atom:
H—Cl
H Cl
1 + 7 = 8 electrons
2. Bond the atoms—sharing 2 electrons. Show all valence electrons for each atom.
..
H—Cl:
¨
3. Count all the electrons in the bonds and electrons that are not bonded (unshared, non-bonded, lone pairs), they
should equal the total number of valence electrons.
..
H—Cl:
¨
2 e– (from bond)
+ 6 unshared electrons around the chloride atom
8 total electrons
Polyatomic compounds that obey the octet rule:
1. Determine the total number of valence electrons
2. Determine central atom
3. Draw a bond for each atom that is bonded to the central atom.
4. Give each atom that requires an octet the appropriate number of electrons.
5. Count all electrons (included bonded atoms)
6. The total number of electrons should equal the number of valence electrons available.
CH4,
1. CH4 has 4+4 = 8 valence electrons
2. Carbon is the central atom
3. Add hydrogen bonds
H
׀
H ̶ C ̶ H
׀
H
4. Count the total number of electrons
NH3
1. NH3 has 5+3 = 8 valence electrons
2. Nitrogen is the central atom
3. Add hydrogen bonds
4. Add two electrons to nitrogen; it must have an octet.
..
H ̶ N ̶ H
׀
H
3. Count the total number of electrons
6. 6
H2O
1. H2O has 5+3 = 8 valence electrons
2. Oxygen is the central atom
3. Add hydrogen bonds
4. Add electrons to oxygen, two at a time, until it has an octet ( it must have an octet).
..
H ̶ O:
׀
H
5. Count the total number of electrons
More examples will be given in class.
Polyatomic compounds that do not obey the octet rule:
BeBr2,
• BeBr ,
2
– has a total of 16 valence electrons
– Be is the central atom
– Give each bromine an octet
.. ..
: Br ̶ Be ̶ Br :
¨ ¨
Count the electrons; all the valence electrons have been used.
BCl3
– has a total of 24 valence electrons
– B is the central atom
– Give each chlorine an octet
..
: Cl :
׀
B
.. ..
: Cl Cl:
¨ ¨
– Count the electrons—all the valence electrons have been used.
Double bonds and triple bonds:
O2
– has a total of 12 valence electrons
– either O is the central atom
– each oxygen must have an octet
– There are not enough valence electrons to give each nitrogen and octet; therefore, we must form double
bonds.
.. ..
:O=O:
7. 7
N2
– has a total of 10 valence electrons
– either N is the central atom
– each nitrogen must have an octet
– There are not enough valence electrons to give each nitrogen and octet; therefore, we must form triple
bonds.
:N≡N:
CO2
– has a total of 16 valence electrons
– either C is the central atom
– form a bond with each oxygen
– each oxygen must have an octet
– There are not enough valence electrons to give each oxygen an octet; therefore, we must form double
bonds.
.. ..
: O =C ̶ O :
¨ ¨
Resonance structures: Delocalized electrons determined by bond strength. Double bonds are shorter and stronger
than single bonds. (Bond strength , the amount of energy required to break a bond)
Write Lewis dot and dash formulas for sulfur trioxide, SO .
3
– 4 x 6 = 24 valence electrons
– sulfur is central atom
– give all oxygen an octet , a double bond will have to be formed between one oxygen and the sulfur
central atom.
·· ··
· O · S· · O ·
· or ·O S O·
· ·· · ·
·· ··
·
·· ·· ··
· O · ·O·
· ·· ·
· ·
··
There are three possible structures for SO3.
– The double bond can be placed in one of three places.
·· ·· ·· ··
·O S O ·
· ·
·O
· S O· · ·O S O ·
· ·
·· ·· ·· ·· ·· ··
·O ·
· ·· · ·O· ·O ·
· ·· ·
· ·
When two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance
structures to show the molecule’s structure. Double-headed arrows are used to indicate resonance formulas.
Resonance is a flawed method of representing molecules.
There are no single or double bonds in SO .
– 3
In fact, all of the bonds in SO3 are equivalent.
The best Lewis formula of SO that can be drawn is:
3
O S O
We will do these in class.
SO2 O
8. 8
NO3
Polyatomic ions: The charges on the ion must be either added or substracted from the total valence electrons.
NH4+ (Must substract an electron because the molecule has a + 1 charge.)
5 + 4 – 1 = 8 valence electrons.
We will do these in class:
SO42-
NO2-
CO32-
Salts with polyatomic ions
Shapes and Polarity of Molecules
(VSEPR—valence shell electron pair repulsion theory)
Electronic shape—counting regions of high electron density
1. Each bond is counted as one region of high electron density.
2. Each double bond and each triple bond is counted as one region of high electron density.
3. Each pair of unshared electrons is counted as one region of high electron density.
9. 9
Rules:
1. Determine the central atom
2. Determine the number of high electron densities around the central atom.
3. Determine the electronic structure based on geometric structure of molecule.
4. Polarity of a molecule is determined by the symmetry of the electron density around the central atom. (The
Electronegativity must be equal and opposite in magnitude)
Number of high Electronic
Molecule Central Atom Polarity
electron densities structure
BeCl2 Be 2 linear nonpolar
BCl3 B 3 trigonal planar nonpolar
CH4 C 4 tetrahedral nonpolar
NH3 N 4 tetrahedral polar
H2O O 4 tetrahedral polar
trigonal
PCl5 P 5 nonpolar
bipyramidal
SF6 S 6 octrahedral nonpolar