This document provides an overview of ionic bonding. Ionic bonds form between elements when one atom loses electrons to become a positively charged ion and another atom gains those electrons to become a negatively charged ion. For example, in sodium chloride, sodium atoms lose electrons to form Na+ ions and chloride atoms gain electrons to form Cl- ions. The oppositely charged ions are then held together by electrostatic forces in a repeating crystal lattice structure. Ionic compounds have high melting points, are brittle, and do not conduct electricity in solid form but do conduct when molten or dissolved in water as the ions become mobile.
The document discusses the different types of bonding mechanisms that hold atoms together in solids, including ionic bonding, covalent bonding, metallic bonding, van der Waals bonding, and hydrogen bonding. Ionic bonding involves the transfer of electrons between metals and nonmetals, covalent bonding involves the sharing of electrons between nonmetals, and metallic bonding involves the delocalization of electrons among metal atoms. Weaker van der Waals forces result from induced dipole interactions between molecules. Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen or fluorine.
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.
Ionic compounds are formed through ionic bonding between metals and nonmetals. Electrons are transferred from the metal atoms to the nonmetal atoms, resulting in cations with positive charges and anions with negative charges. The electrostatic forces between the oppositely charged ions hold the compound together in a crystalline lattice structure. Common polyatomic ions like nitrate, sulfate, and phosphate are also present in ionic compounds. Ionic compounds have properties like being solid at room temperature, having high melting points, and being good conductors of electricity when molten or dissolved in water.
Ionic bond seminar by Mohammad Nasih
in Kurdistan -Iraq
Kurdistan regional government
Ministry of higher education & scientific research
University scientific
Part chemistry
Introduction
Some Information & Properties about Ionic Bonding
Write Chemical Formula about this substance
Atoms gain or lose
Formation of Ions from Metals
Ions from Nonmetal Ions
Some Typical Ions with Positive Charges (Cations)
Ionic bonds form between metal and nonmetal atoms when electrons are transferred. Positively charged metal ions are called cations, and negatively charged nonmetal ions are called anions. Ionic bonds occur when cations and anions are attracted to each other due to their opposite charges, balancing out the overall charge. Ionic compounds write their chemical formulas to represent this charge balance between the ions.
Metallic bonding involves a lattice of positive metal ions surrounded by delocalized electrons that form a negative "electron cloud". This electron cloud binds the positively charged ions together and is responsible for metals' properties. Metals exhibit high electrical conductivity because the mobile electron cloud allows electrons to move freely throughout the structure. The strength of metallic bonding depends on electron density and ionic radius - greater electron density and smaller ionic radius lead to stronger bonding and higher melting points. Alloys are mixtures of metals designed to enhance properties like strength, corrosion resistance, magnetism, and ductility.
This document provides an overview of ionic bonding. Ionic bonds form between elements when one atom loses electrons to become a positively charged ion and another atom gains those electrons to become a negatively charged ion. For example, in sodium chloride, sodium atoms lose electrons to form Na+ ions and chloride atoms gain electrons to form Cl- ions. The oppositely charged ions are then held together by electrostatic forces in a repeating crystal lattice structure. Ionic compounds have high melting points, are brittle, and do not conduct electricity in solid form but do conduct when molten or dissolved in water as the ions become mobile.
The document discusses the different types of bonding mechanisms that hold atoms together in solids, including ionic bonding, covalent bonding, metallic bonding, van der Waals bonding, and hydrogen bonding. Ionic bonding involves the transfer of electrons between metals and nonmetals, covalent bonding involves the sharing of electrons between nonmetals, and metallic bonding involves the delocalization of electrons among metal atoms. Weaker van der Waals forces result from induced dipole interactions between molecules. Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like oxygen or fluorine.
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.
Ionic compounds are formed through ionic bonding between metals and nonmetals. Electrons are transferred from the metal atoms to the nonmetal atoms, resulting in cations with positive charges and anions with negative charges. The electrostatic forces between the oppositely charged ions hold the compound together in a crystalline lattice structure. Common polyatomic ions like nitrate, sulfate, and phosphate are also present in ionic compounds. Ionic compounds have properties like being solid at room temperature, having high melting points, and being good conductors of electricity when molten or dissolved in water.
Ionic bond seminar by Mohammad Nasih
in Kurdistan -Iraq
Kurdistan regional government
Ministry of higher education & scientific research
University scientific
Part chemistry
Introduction
Some Information & Properties about Ionic Bonding
Write Chemical Formula about this substance
Atoms gain or lose
Formation of Ions from Metals
Ions from Nonmetal Ions
Some Typical Ions with Positive Charges (Cations)
Ionic bonds form between metal and nonmetal atoms when electrons are transferred. Positively charged metal ions are called cations, and negatively charged nonmetal ions are called anions. Ionic bonds occur when cations and anions are attracted to each other due to their opposite charges, balancing out the overall charge. Ionic compounds write their chemical formulas to represent this charge balance between the ions.
Metallic bonding involves a lattice of positive metal ions surrounded by delocalized electrons that form a negative "electron cloud". This electron cloud binds the positively charged ions together and is responsible for metals' properties. Metals exhibit high electrical conductivity because the mobile electron cloud allows electrons to move freely throughout the structure. The strength of metallic bonding depends on electron density and ionic radius - greater electron density and smaller ionic radius lead to stronger bonding and higher melting points. Alloys are mixtures of metals designed to enhance properties like strength, corrosion resistance, magnetism, and ductility.
Ionic bonding occurs when atoms of metals and non-metals combine to form ionic compounds. Atoms of metals will donate electrons to form cations, while atoms of non-metals will accept electrons to form anions. This transfer of electrons allows the atoms to achieve stable electron configurations similar to noble gases. Common examples are sodium chloride, which forms when sodium donates an electron to chlorine, and magnesium oxide, which forms when magnesium donates two electrons to oxygen. The chemical formulas of ionic compounds are written to balance the charges of the cation and anion.
This document discusses three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds.
1. Ionic bonds form between metals and nonmetals by the transfer of electrons from one atom to another, resulting in positively and negatively charged ions.
2. Covalent bonds form between nonmetals of similar electronegativity by the sharing of electron pairs between atoms.
3. Metallic bonds form between metallic elements and involve a sea of electrons that hold the metal atoms together strongly.
Definition of ionic compounds
Property of ionic compoundIn crystal formSolubility in waterHigh melting and boiling pointConductivity of electricity
Use of ionic compoundUse as dryersUse to make pans
1) Atoms are made up of electrons, protons and neutrons. Electrons occupy electron shells and determine how an atom bonds with other atoms.
2) Ionic bonds form when a metal transfers electrons to a nonmetal, giving them opposing charges. For example, sodium (Na) transfers an electron to chlorine (Cl), forming sodium ion (Na+) and chloride ion (Cl-).
3) Covalent bonds form when atoms share electrons to fill their outer electron shells in a stable configuration.
Metallic bonding occurs when metal atoms lose valence electrons to become positively charged ions embedded in a "sea" of delocalized electrons. This electron sea model explains several properties of metals, including their ability to conduct electricity and heat, as well as their malleability, ductility, lustrous appearance, and high melting and boiling points. The mobile electrons allow for heat and charge conduction, while the metallic lattice structure enables atoms to slide past one another under pressure.
The document discusses electronegativity and types of chemical bonding. Electronegativity is the ability of an atom to attract shared electrons, and the difference in electronegativity between two atoms determines the type of bond formed. There are three main types of bonds - covalent, ionic, and metallic - which give compounds different physical properties such as state, melting/boiling points, conductivity, and solubility. The document also lists topics to cover in more depth related to chemical bonding concepts.
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.
Ionic bonds form between a metal cation and nonmetal anion when the metal loses electrons to the nonmetal to achieve a stable electron configuration. Covalent bonds involve the sharing of electron pairs between two nonmetal atoms and result in weaker bonds than ionic. Metallic bonds are formed by the attraction between free-floating valence electrons in metals and the fixed positively charged metal ions.
This document discusses ionic compounds and the formation of ionic bonds. Ions form when atoms gain or lose electrons to achieve stable electron configurations like noble gases. Ionic bonds occur between oppositely charged ions and result in crystalline solids with high melting points. The document explains how to name ionic compounds based on the cation and anion present and write chemical formulas from compound names. It also briefly discusses metallic bonding and the properties of metals and alloys.
The document provides an overview of chemical bonding principles including ionic bonding, covalent bonding, and molecular geometry. It discusses how atoms interact to achieve stable electronic configurations through ionic bonding by exchanging electrons or covalent bonding by sharing electrons. Lewis structures are introduced as a way to represent valence electrons in molecules and determine molecular geometry and polarity based on electron pair arrangements around central atoms. Key concepts covered include octet rule, electronegativity, bond polarity, and using Lewis structures to systematically determine molecular structure characteristics.
This document summarizes key concepts about valence electrons, electron dot diagrams, ionic bonding, and covalent bonding. Valence electrons are found in the outermost energy level and interact during bonding. Electron dot diagrams represent valence electrons with dots. Ionic bonding occurs through the transfer of valence electrons between metals and nonmetals, resulting in ions with charges. Covalent bonding involves the sharing of valence electrons between nonmetals. Properties of ionic bonds include conductivity, higher melting/boiling points, and crystal formation, while covalent substances are often gases/liquids and nonconductive.
This document discusses different types of bonds:
- Ionic bonds form when a metal transfers valence electrons to a nonmetal, resulting in positively charged metal ions and negatively charged nonmetal ions that are attracted to each other. Ionic compounds are usually solid and brittle with high melting points.
- Covalent bonds form when nonmetal atoms share valence electrons to form molecules. Covalent compounds consist of many molecules.
- Metallic bonds form when metal atoms share their valence electrons, resulting in a "sea of electrons" that surrounds the positive metal ions and allows metals to conduct electricity and heat well.
The document discusses two types of bonding: covalent and ionic. Covalent bonding involves the sharing of electron pairs between non-metal atoms to form molecules. Ionic bonding occurs between a metal and non-metal where the metal transfers electrons to the non-metal to form oppositely charged ions in a giant lattice structure. This ionic lattice structure results in high melting and boiling points due to the strong electrostatic forces between the ions.
There are three main types of chemical bonds:
1) Ionic bonds form when electrons are transferred from a metal to a non-metal atom. They result in charged ions that are attracted to each other.
2) Covalent bonds form when atoms share electrons through single, double or triple bonds. Covalent compounds have low melting points and are electrical insulators.
3) Dative bonds form when one atom donates a lone pair of electrons to an atom with an incomplete octet.
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.
The document discusses different types of chemical bonds including ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons between metals and non-metals, resulting in ionic compounds with high melting and boiling points that conduct electricity when melted or in solution. Covalent bonds involve the sharing of electrons between non-metal atoms, resulting in molecules with low melting and boiling points that do not conduct electricity. The document also discusses metallic bonding and how metal atoms are held together by valence electrons that are delocalized throughout the structure.
This document discusses the nature of ionic bonds and properties of ionic compounds. It explains that ionic bonding occurs when there is a large difference in electronegativity (>1.7) between a metal and non-metal atom. The metal atom donates electrons to become a positively charged cation, while the non-metal atom gains electrons to become a negatively charged anion. These oppositely charged ions are arranged in a regular crystal lattice structure held together by electrostatic forces. Ionic compounds have properties stemming from this lattice structure such as being crystalline solids, brittle, and able to conduct electricity when in solution but not in solid form.
The document discusses different types of chemical bonds:
- Ionic bonds form when there is a large difference in electronegativity between two atoms, causing one atom to transfer an electron to the other.
- Covalent bonds form when there is a small difference in electronegativity between two non-metal atoms, causing them to share electrons. Covalent bonds can be nonpolar or polar.
- Water is an example of a polar covalent molecule, with the oxygen end being partially negative and the hydrogen ends being partially positive.
Hello everyone, I am Dr. Ujwalkumar Trivedi, Head of Biotechnology Department at Marwadi University Rajkot. I teach Molecular Biology to the students of M.Sc. Microbiology and Biotechnology.
The current presentation talks about the formation of chemical bonds. This presentation gives insight into the formation of Ionic Bonds, Covalent Bonds and Metallic Bonds with examples.
1. Hess's law states that the overall enthalpy change of a chemical process is independent of the path taken.
2. The document uses the reaction of sodium hydroxide and hydrochloric acid to illustrate Hess's law. It shows that the enthalpy change of the direct reaction is equal to the sum of the enthalpy changes for the stepwise reactions.
3. Representing the reactions on an enthalpy level diagram and enthalpy cycle confirms that Hess's law applies, with the enthalpy change for the overall reaction being the sum of the enthalpy changes for the individual steps.
Ionic compounds are formed from a cation and an anion. Opposite charges attract the ions together in an ionic bond. Ions transfer electrons to achieve stable noble gas electron configurations. Ionic compounds have crystalline structures where ions are arranged in repeating patterns. The ions are strongly bonded and give ionic solids properties like brittleness and ability to conduct electricity when molten or dissolved in water.
Ionic bonding occurs when atoms of metals and non-metals combine to form ionic compounds. Atoms of metals will donate electrons to form cations, while atoms of non-metals will accept electrons to form anions. This transfer of electrons allows the atoms to achieve stable electron configurations similar to noble gases. Common examples are sodium chloride, which forms when sodium donates an electron to chlorine, and magnesium oxide, which forms when magnesium donates two electrons to oxygen. The chemical formulas of ionic compounds are written to balance the charges of the cation and anion.
This document discusses three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds.
1. Ionic bonds form between metals and nonmetals by the transfer of electrons from one atom to another, resulting in positively and negatively charged ions.
2. Covalent bonds form between nonmetals of similar electronegativity by the sharing of electron pairs between atoms.
3. Metallic bonds form between metallic elements and involve a sea of electrons that hold the metal atoms together strongly.
Definition of ionic compounds
Property of ionic compoundIn crystal formSolubility in waterHigh melting and boiling pointConductivity of electricity
Use of ionic compoundUse as dryersUse to make pans
1) Atoms are made up of electrons, protons and neutrons. Electrons occupy electron shells and determine how an atom bonds with other atoms.
2) Ionic bonds form when a metal transfers electrons to a nonmetal, giving them opposing charges. For example, sodium (Na) transfers an electron to chlorine (Cl), forming sodium ion (Na+) and chloride ion (Cl-).
3) Covalent bonds form when atoms share electrons to fill their outer electron shells in a stable configuration.
Metallic bonding occurs when metal atoms lose valence electrons to become positively charged ions embedded in a "sea" of delocalized electrons. This electron sea model explains several properties of metals, including their ability to conduct electricity and heat, as well as their malleability, ductility, lustrous appearance, and high melting and boiling points. The mobile electrons allow for heat and charge conduction, while the metallic lattice structure enables atoms to slide past one another under pressure.
The document discusses electronegativity and types of chemical bonding. Electronegativity is the ability of an atom to attract shared electrons, and the difference in electronegativity between two atoms determines the type of bond formed. There are three main types of bonds - covalent, ionic, and metallic - which give compounds different physical properties such as state, melting/boiling points, conductivity, and solubility. The document also lists topics to cover in more depth related to chemical bonding concepts.
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.
Ionic bonds form between a metal cation and nonmetal anion when the metal loses electrons to the nonmetal to achieve a stable electron configuration. Covalent bonds involve the sharing of electron pairs between two nonmetal atoms and result in weaker bonds than ionic. Metallic bonds are formed by the attraction between free-floating valence electrons in metals and the fixed positively charged metal ions.
This document discusses ionic compounds and the formation of ionic bonds. Ions form when atoms gain or lose electrons to achieve stable electron configurations like noble gases. Ionic bonds occur between oppositely charged ions and result in crystalline solids with high melting points. The document explains how to name ionic compounds based on the cation and anion present and write chemical formulas from compound names. It also briefly discusses metallic bonding and the properties of metals and alloys.
The document provides an overview of chemical bonding principles including ionic bonding, covalent bonding, and molecular geometry. It discusses how atoms interact to achieve stable electronic configurations through ionic bonding by exchanging electrons or covalent bonding by sharing electrons. Lewis structures are introduced as a way to represent valence electrons in molecules and determine molecular geometry and polarity based on electron pair arrangements around central atoms. Key concepts covered include octet rule, electronegativity, bond polarity, and using Lewis structures to systematically determine molecular structure characteristics.
This document summarizes key concepts about valence electrons, electron dot diagrams, ionic bonding, and covalent bonding. Valence electrons are found in the outermost energy level and interact during bonding. Electron dot diagrams represent valence electrons with dots. Ionic bonding occurs through the transfer of valence electrons between metals and nonmetals, resulting in ions with charges. Covalent bonding involves the sharing of valence electrons between nonmetals. Properties of ionic bonds include conductivity, higher melting/boiling points, and crystal formation, while covalent substances are often gases/liquids and nonconductive.
This document discusses different types of bonds:
- Ionic bonds form when a metal transfers valence electrons to a nonmetal, resulting in positively charged metal ions and negatively charged nonmetal ions that are attracted to each other. Ionic compounds are usually solid and brittle with high melting points.
- Covalent bonds form when nonmetal atoms share valence electrons to form molecules. Covalent compounds consist of many molecules.
- Metallic bonds form when metal atoms share their valence electrons, resulting in a "sea of electrons" that surrounds the positive metal ions and allows metals to conduct electricity and heat well.
The document discusses two types of bonding: covalent and ionic. Covalent bonding involves the sharing of electron pairs between non-metal atoms to form molecules. Ionic bonding occurs between a metal and non-metal where the metal transfers electrons to the non-metal to form oppositely charged ions in a giant lattice structure. This ionic lattice structure results in high melting and boiling points due to the strong electrostatic forces between the ions.
There are three main types of chemical bonds:
1) Ionic bonds form when electrons are transferred from a metal to a non-metal atom. They result in charged ions that are attracted to each other.
2) Covalent bonds form when atoms share electrons through single, double or triple bonds. Covalent compounds have low melting points and are electrical insulators.
3) Dative bonds form when one atom donates a lone pair of electrons to an atom with an incomplete octet.
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.
The document discusses different types of chemical bonds including ionic bonds and covalent bonds. Ionic bonds involve the transfer of electrons between metals and non-metals, resulting in ionic compounds with high melting and boiling points that conduct electricity when melted or in solution. Covalent bonds involve the sharing of electrons between non-metal atoms, resulting in molecules with low melting and boiling points that do not conduct electricity. The document also discusses metallic bonding and how metal atoms are held together by valence electrons that are delocalized throughout the structure.
This document discusses the nature of ionic bonds and properties of ionic compounds. It explains that ionic bonding occurs when there is a large difference in electronegativity (>1.7) between a metal and non-metal atom. The metal atom donates electrons to become a positively charged cation, while the non-metal atom gains electrons to become a negatively charged anion. These oppositely charged ions are arranged in a regular crystal lattice structure held together by electrostatic forces. Ionic compounds have properties stemming from this lattice structure such as being crystalline solids, brittle, and able to conduct electricity when in solution but not in solid form.
The document discusses different types of chemical bonds:
- Ionic bonds form when there is a large difference in electronegativity between two atoms, causing one atom to transfer an electron to the other.
- Covalent bonds form when there is a small difference in electronegativity between two non-metal atoms, causing them to share electrons. Covalent bonds can be nonpolar or polar.
- Water is an example of a polar covalent molecule, with the oxygen end being partially negative and the hydrogen ends being partially positive.
Hello everyone, I am Dr. Ujwalkumar Trivedi, Head of Biotechnology Department at Marwadi University Rajkot. I teach Molecular Biology to the students of M.Sc. Microbiology and Biotechnology.
The current presentation talks about the formation of chemical bonds. This presentation gives insight into the formation of Ionic Bonds, Covalent Bonds and Metallic Bonds with examples.
1. Hess's law states that the overall enthalpy change of a chemical process is independent of the path taken.
2. The document uses the reaction of sodium hydroxide and hydrochloric acid to illustrate Hess's law. It shows that the enthalpy change of the direct reaction is equal to the sum of the enthalpy changes for the stepwise reactions.
3. Representing the reactions on an enthalpy level diagram and enthalpy cycle confirms that Hess's law applies, with the enthalpy change for the overall reaction being the sum of the enthalpy changes for the individual steps.
Ionic compounds are formed from a cation and an anion. Opposite charges attract the ions together in an ionic bond. Ions transfer electrons to achieve stable noble gas electron configurations. Ionic compounds have crystalline structures where ions are arranged in repeating patterns. The ions are strongly bonded and give ionic solids properties like brittleness and ability to conduct electricity when molten or dissolved in water.
This document discusses ionic and metallic bonding. It explains that ions are formed when atoms gain or lose electrons to achieve stable noble gas electron configurations. Metals form cations by losing electrons while nonmetals form anions by gaining electrons. Ionic compounds contain cations and anions in ratios represented by chemical formulas. Metallic bonding occurs via delocalized valence electrons that are shared between metal atoms.
Chapter 6.3 : Ionic Bonding and Ionic CompoundsChris Foltz
Ionic compounds are composed of positive and negative ions in ratios that balance the charges. They form crystalline lattices with ions arranged in an orderly pattern. The ions are held together by strong electrostatic forces called lattice energy. Ionic compounds have high melting and boiling points, are hard and brittle, and conduct electricity when molten or dissolved but not in the solid state. Polyatomic ions are charged groups of covalently bonded atoms that behave as single ions in ionic compounds.
Ionic bonds form when a metal transfers an electron to a nonmetal, giving each atom an octet of electrons. For example, sodium loses an electron to form Na+ while chlorine gains that electron to form Cl-. The resulting ions are held together by electrostatic attraction to form an ionic compound, sodium chloride (NaCl). NaCl crystallizes into a repeating pattern where Na+ and Cl- ions alternate in a crystal lattice. Ionic compounds conduct electricity when molten or dissolved due to the movement of ions.
Here are the key points about polar and nonpolar covalent bonds:
- Nonpolar covalent bonds form between atoms with similar electronegativity. Electrons are equally shared. Examples include H2, Cl2, etc.
- Polar covalent bonds form between atoms with different electronegativity. Electrons are unequally shared, resulting in partial positive and negative charges on the atoms. Examples include HCl, H2O.
- Polar molecules are attracted to each other due to their partial charges. They are soluble in polar solvents like water. Nonpolar molecules are not attracted and are soluble in nonpolar solvents like hexane.
- Most bonds have some ionic character based on electrone
A covalent bond forms when two atoms share electrons between their outer shells to satisfy the Octet Rule, which requires atoms to gain, lose, or share electrons until each has eight in their outer valence shell, matching the stable electron configuration of noble gases. For example, phosphorus shares electrons with other atoms because it has only five in its outer shell, while gaining three more allows it to match the full outer shell of the noble gas argon.
The Use of Montomorillonte as an absorbent for ignitable liquids from porcine...Matthew Perryman
This document is a thesis submitted in partial fulfillment of a Bachelor of Science degree from Anglia Ruskin University. It examines the use of montmorillonite clay as an absorbent for ignitable liquids from porcine skin. The thesis acknowledges those who helped and supported the author in completing the work. It then provides a table of contents outlining the various sections of the thesis, including an introduction discussing fire chemistry and investigation, adsorption and montmorillonite clay, methods used, and aims of the research.
1. The document discusses how to determine the shape of covalent molecules by counting electrons and bonded pairs to identify the molecular geometry.
2. Regular covalent molecules follow set geometries based on the number of bonded pairs, such as linear for 2 pairs and tetrahedral for 4 pairs.
3. Irregular molecules have lone electron pairs that cause bond angles to decrease from the regular geometry by approximately 2.5 degrees per lone pair.
The document discusses chemical bonding and Lewis structures. It begins by defining a chemical bond as the force that holds atoms together, and discusses how atoms combine or share electrons to form ionic or covalent bonds. It then explains Lewis structures, showing how to draw the Lewis dot symbols and determine the hybridization of atoms. Examples are provided of writing Lewis structures for different molecules like CCl4 and NH4+. The document also discusses exceptions to the octet rule and theories of covalent bonding like valence bond theory and hybridization theory.
The document discusses valence electrons and bonding. It introduces the 2-8-8 rule which states that the first energy level can hold up to 2 electrons, and the second and third levels can each hold up to 8 electrons. Atoms are stable once their energy levels are filled. Ionic bonds form when a metal atom transfers electrons to a nonmetal, becoming cations and anions. Covalent bonds form when atoms share electrons rather than transfer them. Metallic bonds form a "sea" of delocalized electrons between positively charged metal ions.
Covalent bonds result from the sharing of valence electrons between two nonmetallic atoms. A single covalent bond forms when one pair of electrons is shared, while double and triple bonds share two or three pairs of electrons, respectively. Molecular structure and bond angles are determined by the VSEPR model, which predicts molecular geometry based on electron pair repulsion. Hybridization occurs when atomic orbitals mix to form new hybrid orbitals and explain molecular bonding orientations.
1) Noble gases are unreactive because their outer electron shells are full, giving them a stable configuration.
2) Atoms form ions to achieve a full outer shell like noble gases. Metals form cations by losing electrons, and non-metals form anions by gaining electrons.
3) Ionic bonds form when oppositely charged ions are attracted to each other, such as sodium and chloride ions forming sodium chloride. The chemical formula shows the ratio of ions.
Ionic compounds form crystals and have high melting and boiling points, making them very hard and brittle. They are good insulators but conduct electricity when dissolved in water. Common ionic bonds occur between metals like sodium and nonmetals like chlorine, as seen in sodium chloride.
Ionic bonding occurs when atoms transfer electrons to form ions with opposite charges that are attracted via electrostatic forces. Metals form cations by losing electrons to achieve stable electron configurations like noble gases, while nonmetals form anions by gaining electrons. This transfer of electrons allows the formation of ionic compounds with crystalline structures where ion attractions are maximized and repulsions minimized. Properties of ionic compounds include high melting points, solubility in water, defined crystal structures, and the ability to conduct electricity when molten. Metallic bonding also involves cations but is characterized by delocalized valence electrons that form a "sea" allowing metals to conduct electricity and be malleable and ductile.
The document summarizes key concepts about covalent bonds, including how they form through the sharing of valence electrons between atoms to achieve stability. It discusses ionic vs. covalent bonds, how covalent bonds are represented using Lewis dot structures, and how the octet rule determines bonding patterns for atoms in different groups on the periodic table. Examples of writing Lewis dot structures for several diatomic and polyatomic molecules are provided.
It's very good for SPM students . You have to learn the ionic bond thoroughly. If you understand well you can explain it vividly. For other chemistry notes can email me puterizamrud@gmail.com or facebook Pusat Tuisyen Zamrud .
Ionic bonding occurs when atoms of metals and non-metals combine to form ionic compounds. Atoms of metals will donate electrons to form cations, while atoms of non-metals will accept electrons to form anions. This transfer of electrons allows the atoms to achieve stable electron configurations similar to noble gases. Common examples are sodium chloride, which forms when sodium donates an electron to chlorine, and magnesium oxide, which forms when magnesium donates two electrons to oxygen. The chemical formulas of ionic compounds are written to balance the charges of the cation and anion.
This document discusses the three main types of chemical bonds: ionic, covalent, and metallic. Ionic bonds form between metals and nonmetals when the metal transfers electrons to the nonmetal. Covalent bonds form when atoms share electrons in either single, double or triple bonds. Metallic bonds form when metal atoms contribute electrons to form a "sea of electrons" that are shared between all the atoms.
2016 topic 5.1 measuring energy changesDavid Young
This document provides an overview of exothermic and endothermic reactions, calorimetry, and how to calculate enthalpy changes. The key points are:
- Exothermic reactions release heat while endothermic reactions absorb heat.
- Enthalpy change (ΔH) is the quantity of heat released or absorbed during a chemical reaction.
- Calorimetry experiments allow calculation of ΔH by measuring the temperature change of a reaction mixture.
- Sample problems demonstrate how to use calorimetry data like mass changes and temperature differences to calculate the enthalpy change for a reaction.
A projectile is an object moving under the influence of gravity with a parabolic path. Its motion can be analyzed by separating the horizontal and vertical components. In the horizontal direction, the velocity is constant, while in the vertical direction there is a constant acceleration due to gravity. Projectile motion is used to model many real-world scenarios like thrown objects, diving, and artillery fire. Solving projectile motion problems involves separating the horizontal and vertical motions and using kinematic equations with the initial velocities and gravitational acceleration.
1) Acceleration is a measure of how quickly velocity changes. It can represent a change in speed, direction, or both.
2) The slope of a velocity-time graph represents acceleration. The area under the graph represents displacement.
3) For objects experiencing uniform acceleration, displacement can be calculated using: s = ut + (1/2)at^2
1.2 displacement and position vs time graphsDavid Young
1) Displacement is defined as the straight line path between an object's initial and final positions, representing the change in position. It is not necessarily the same as the total distance travelled.
2) Victor's displacement is calculated as his final position (2km North of Starbucks) minus his initial position (3km South of Starbucks), equaling 5km North.
3) A position-time graph shows an object's position over time. The slope of the line between any two points represents the average velocity during that time interval.
1) Vectors are quantities that have both magnitude and direction, unlike scalars which only have magnitude. Common vector quantities include force, velocity, and displacement.
2) Vectors can be described by their magnitude, direction, and reference frame. Direction is measured in degrees from the x-axis.
3) Vector operations like addition, subtraction, multiplication and division can be performed graphically or using components and trigonometry. Parallel vectors can be added/subtracted algebraically but others require using components.
4) To add vectors using components, each vector is broken into x and y components using trigonometry. The x and y components can then be added separately and combined to find the overall resultant vector.
This document discusses the basics of sound waves. It explains that sound waves are caused by vibrations, requiring a medium like air, liquid or solid to transfer energy. It describes how compression and rarefaction of molecules transfers sound waves. Frequency determines pitch, with higher frequencies being higher pitched. Amplitude determines loudness, with greater amplitude being louder. Graphs can show the displacement over time of sound waves to represent characteristics like frequency and amplitude. The speed of sound depends on the medium and temperature.
The document summarizes key concepts about waves, including:
1) Waves can be classified as mechanical or electromagnetic depending on whether they require a medium to travel. Mechanical waves include sound and water waves while electromagnetic waves include radio and light waves.
2) Waves can be transverse, with oscillations perpendicular to the direction of travel, or longitudinal, with oscillations parallel to the direction of travel.
3) Key wave properties include frequency, wavelength, period, amplitude, and speed. The speed of a wave depends on properties of the medium and can be calculated from the wavelength and period.
This document discusses quantitative chemistry concepts including the mole, Avogadro's constant, molar mass, and using moles to calculate the number of particles and mass of substances. Some key points covered are:
- A mole is a unit used to count fundamental particles and represents 6.02x1023 particles.
- The mole allows converting between large numbers of particles and more manageable units like grams.
- Molar mass relates the mass of a sample of a substance to the number of moles it contains.
- Problems use molar mass to calculate moles, particles, or mass given any two of those quantities.
- Atoms consist of a nucleus containing protons and neutrons surrounded by electrons in orbitals.
- The atomic number is the number of protons, which identifies the element. The mass number is the total number of protons and neutrons.
- Isotopes are atoms of the same element with different numbers of neutrons. The relative atomic mass takes into account the natural abundance of isotopes.
This document discusses electron configuration and how it is described using quantum numbers. It explains the main concepts like energy shells (n), subshells (s, p, d, f), orbitals and how electrons fill these according to the Aufbau principle and Hund's rule. Examples are provided to show the electron configurations of elements like hydrogen, helium, lithium and how the configurations change as the atomic number increases. Practice problems are included at the end to determine configurations of additional elements.
This document discusses atomic structure and properties. It begins by reviewing the basic atomic model including protons, neutrons, and electrons. It then discusses atomic number, mass number, isotopes, ions, and electron configuration. The document also covers average atomic mass and mass spectrometry. It concludes by discussing the Bohr model of the atom and how this helped explain emission spectra of elements. In summary, the key topics covered are the fundamental particles of atoms, atomic notation, isotopes, ions, electron configuration, and early atomic structure models.
2016 topic 0 - oxidation and reduction (INTRODUCTION)David Young
The document discusses oxidation and reduction reactions. It defines oxidation as the loss of electrons or an increase in oxidation state, and reduction as the gain of electrons or a decrease in oxidation state. Redox reactions involve both oxidation and reduction occurring simultaneously. Oxidation states can be used to identify what is oxidized and reduced in a reaction. The document provides examples of calculating oxidation states in various compounds and ions using rules like the sum of oxidation states equaling the overall charge.
2016 topic 0 - elements & periodic tableDavid Young
- An element is a substance composed of a single type of atom that cannot be broken down further. A compound is a substance made of two or more elements chemically bonded together.
- Elements can exist as individual atoms, molecules of two bonded atoms, or giant molecular or network structures made of many bonded atoms like metals. Elements are categorized as metals, nonmetals, or metalloids based on their physical and chemical properties.
Introduction to elements & periodic tableDavid Young
- An element is a substance composed of a single type of atom that cannot be broken down further. A compound is a substance made of two or more elements chemically bonded together.
- Elements can exist as individual atoms, molecules of two bonded atoms, or giant molecular or network structures made of many bonded atoms like metals. Elements are categorized as metals, nonmetals, or metalloids based on their physical and chemical properties.
Physics ii djy 2013 ppt wave characteristicsDavid Young
This document discusses key characteristics and concepts related to waves, including:
- Waves are patterns of disturbances caused by the movement of energy through matter or space. They can be classified as mechanical or electromagnetic waves.
- Key wave characteristics include frequency, wavelength, amplitude, and speed. Frequency is the number of waves per second, wavelength is the distance between identical points on waves, and amplitude is the maximum displacement from equilibrium.
- Waves transfer energy, not matter. Particle motion can be transverse (perpendicular to the direction of energy transfer) or longitudinal (parallel to the direction of energy transfer).
- Graphs can be used to describe wave characteristics like displacement over position or time. Speed can be
This document provides an overview of organic chemistry topics related to stereoisomers, including:
1) It defines structural isomers and stereoisomers as different types of isomers, and provides examples to illustrate each type.
2) It discusses the importance of stereochemistry in fields like biochemistry, medicine, and studying reaction mechanisms.
3) It explains key concepts around chiral centers, configurations, enantiomers, and how these relate to a molecule's ability to rotate plane-polarized light and be optically active.
This document discusses organic chemistry topics including condensation reactions which combine two smaller molecules to form one larger molecule with the elimination of a small molecule like water. It provides examples of esterification reactions forming esters and amide formation, as well as condensation polymerization reactions that form polymers from monomers with functional groups undergoing condensation. Specific examples discussed include the formation of terylene, nylon 6,6, and polypeptides from amino acid monomers.
This document discusses elimination reactions where a small molecule is removed from a reactant. It describes the elimination of HBr from bromoalkanes using a dilute NaOH solution, which can cause either a substitution or elimination reaction depending on conditions. The two types of elimination reactions are E2, a bimolecular process without intermediates typical of primary/secondary halides, and E1, a unimolecular reaction with a carbocation intermediate typical of tertiary halides. Hydroxide acts as a base by accepting a proton from the alkyl halide, initiating electron movement that forms a C=C double bond and removes the halide.
This document discusses organic chemistry substitution reactions. It provides information on SN2 nucleophilic substitution reactions including:
1) SN2 reactions occur when a bromoalkane is added to a dilute NaOH solution or reacted with cyanide ions or ammonia.
2) The products of SN2 reactions with cyanide ions and ammonia are nitriles and amines, respectively.
3) The electron density and electronegativity of the nucleophile affects the rate of substitution, with more reactive nucleophiles like CN- having faster reaction rates than less reactive ones like NH3.
4) The leaving ability of the halide also impacts the rate, with iodine
Topic 10 & 20 reaction summary and reviewDavid Young
This document summarizes key organic chemistry reactions for IB Chemistry at both the SL and HL levels. It outlines addition, substitution, elimination, oxidation, reduction, condensation, and polymerization reactions. Specific mechanisms that must be known include homolytic fission, SN1 and SN2 nucleophilic substitution, and elimination reactions. Condensation reactions forming esters and amides are also included.
3. Definition a shared pair of electrons generally with one electron being
supplied by each atom either side of the bond.
atoms are held together
because their nuclei which
have an overall positive charge
are attracted to the shared electrons
Formation between atoms of the same element N2, O2, diamond,
graphite
between atoms of different elements CO2, SO2
on the RHS of the table;
BETWEEN NON-METALS
when one of the elements is in CCl4, SiCl4
the middle of the table;
with head-of-the-group elements BeCl2
with high ionisation energies;
COVALENT BONDING
+ +
4. HYDROGEN
Another hydrogen atom
also needs one electron to
complete its outer shell
Hydrogen atom needs
one electron to
complete its outer shell
H
H
5. HYDROGEN
H H
atoms share a pair of electrons to
form a single covalent bond
A hydrogen MOLECULE is formed
H H
H
H
WAYS TO REPRESENT A MOLECULE of H2
6. METHANE
C
Each hydrogen
atom needs 1
electron to
complete its
outer shell
A carbon atom needs 4
electrons to complete
its outer shell
H
H
H
H
8. METHANE
Carbon shares all 4 of
its electrons to form 4
single covalent bonds
H C H
H
H
H C H
H
H
WAYS TO REPRESENT
THE MOLECULE
9. AMMONIA
N
Each hydrogen
atom needs
one electron to
complete its
outer shell
Nitrogen atom needs 3 electrons
to complete its outer shell
Nitrogen can only share 3 of its
5 electrons otherwise it will
exceed the maximum of 8
A LONE PAIR REMAINS
H
H
H
H N H
H
H N H
H
WAYS TO REPRESENT
THE MOLECULE
10. WATER
O
Each hydrogen
atom needs
one electron to
complete its
outer shell
Oxygen atom needs 2 electrons
to complete its outer shell
Oxygen can only share 2 of its 6
electrons otherwise it will
exceed the maximum of 8
2 LONE PAIRS REMAIN
H
H
H O
H
H O
H
WAYS TO REPRESENT
THE MOLECULE
11. OXYGEN
O
each atom needs two electrons
to complete its outer shell
each oxygen shares 2 of its
electrons to form a
DOUBLE COVALENT BOND
O O O
and NITROGEN
12. Orbital theory – bond length and bond strength
Covalent bonds are formed when orbitals, each containing one electron,
overlap.
The greater the overlap the stronger the bond.
The greater the overlap, the shorter the bond length.
less overlapping orbitals Single < Double < Triple more overlapping orbitals
orbital
containing 1
electron
orbital
containing 1
electron
bonding orbital shape
13. COMPARE Bond lengths and strengths
(data booklet has all the values)
O – O 0.148 146
O = O 0.121 496
C – C 0.154 348
C = C 0.134 612
C ≡ C 0.120 837
C – O 0.143 360
C = O 0.122 743
Multiple Bonds
14. Bond Polarity
A : B
Consider a covalent bond between A and B. If A and B
have the same electronegativity, the electrons are
shared equally.
If B is more electronegative than A, the shared
electrons have a greater probability of being found
closer to B. POLAR COVALENT
+ A : B -
15. HYDROGEN CHLORIDE
Cl H
Hydrogen atom also
needs one electron
to complete its outer
shell
Chlorine atom
needs one electron
to complete its
outer shell
16. HYDROGEN CHLORIDE
Cl H
atoms share a pair of
electrons to form a
single covalent bond
H Cl
H Cl
WAYS TO REPRESENT THE MOLECULE
17. Bonding Atoms are joined together within the molecule by covalent bonds.
Electrical Don’t conduct electricity as they have no mobile ions or electrons
Solubility Tend to be more soluble in organic solvents than in water;
some are hydrolysed
Boiling point Low - compared to ionic compounds
intermolecular forces are weak;
they increase as molecules get a larger surface area
e.g. CH4 -161°C C2H6 - 88°C C3H8 -42°C
as the intermolecular forces are weak, little energy is required to
to separate molecules from each other so boiling points are low
some boiling points are higher than expected for a given mass
because you can get additional forces of attraction
these concepts will be clarified and expanded on in later lessons
SIMPLE COVALENT MOLECULES
19. • atoms share electrons to get the nearest noble gas electronic configuration
NOT ALWAYS
• some don’t achieve an “octet” as they haven’t got enough electrons
ex Al in AlCl3 (why is this a covalent bond?)
• others share only some - if they share all they will exceed their “octet”
ex NH3 and H2O
• atoms of elements in the 3rd period onwards can exceed their “octet” if
they wish as they are not restricted to eight electrons in their “outer shell”
ex PCl5 and SF6 (HL only)
COVALENT BONDING - EXCEPTIONS