Created by Turbolearn AI
Atomic Combinations
Chemical Bonds
Chemical bonds are forces that hold atoms together to form molecules or compounds. These bonds arise from
the interactions between the positive nuclei and the negative electrons of atoms.
Atoms combine to achieve a more stable electron configuration, typically resembling that of a noble gas (8
valence electrons).
Types of Chemical Bonds
1. Covalent Bonds:
Formed by the sharing of electrons between two atoms.
Typically occur between two nonmetal atoms.
Examples: water (H₂O), methane (CH₄)
2. Ionic Bonds:
Formed by the transfer of electrons from one atom to another, resulting in the formation of ions.
Occur between a metal and a nonmetal.
The metal atom loses electrons to become a positively charged ion (cation), and the nonmetal atom gains
electrons to become a negatively charged ion (anion).
These oppositely charged ions are attracted to each other, forming the ionic bond.
Examples: sodium chloride (NaCl), magnesium oxide (MgO)
3. Metallic Bonds:
Found in metals.
Characterized by a "sea" of delocalized electrons that are free to move throughout the metallic structure.
This electron mobility accounts for the high electrical and thermal conductivity of metals.
Examples: copper (Cu), iron (Fe)
Molecular Shape
The shape of a molecule is determined by the arrangement of its atoms in three-dimensional space.
The shape influences a molecule's physical and chemical properties, such as its polarity, reactivity, and
biological activity.
Valence Shell Electron Pair Repulsion (VSEPR) theory helps predict the shape of molecules based on the
repulsion between electron pairs surrounding a central atom.
Basic Molecular Shapes:
Linear: Atoms are arranged in a straight line (e.g., CO₂)
Trigonal Planar: Three atoms are bonded to a central atom, forming a flat, triangular shape (e.g., BF₃)
Tetrahedral: Four atoms are bonded to a central atom, forming a three-dimensional pyramid shape (e.g., CH₄)
Bent: Similar to trigonal planar or tetrahedral but with one or more lone pairs of electrons, resulting in a bent
shape (e.g., H₂O)
Trigonal Pyramidal: Three atoms and one lone pair are bonded to a central atom, forming a pyramid shape (e.g.,
NH₃)
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Electronegativity
Electronegativity is a measure of the ability of an atom in a chemical compound to attract shared
electrons in a covalent bond.
Electronegativity differences between atoms in a bond determine the polarity of the bond.
Linus Pauling scale is commonly used to represent electronegativity.
Electronegativity and Bond Polarity
Nonpolar Covalent Bond:
Occurs when two atoms have similar electronegativities.
Electrons are shared equally.
Example: H₂
Polar Covalent Bond:
Occurs when two atoms have different electronegativities.
Electrons are shared unequally, resulting in a partial positive charge (δ+) on the less electronegative atom
and a partial negative charge (δ−) on the more electronegative atom.
Example: H₂O
Ionic Bond:
Occurs when there is a large difference in electronegativity between two atoms.
Electrons are essentially transferred from one atom to another, resulting in the formation of ions.
Example: NaCl
Energy and Bonding
Energy is involved in the formation and breaking of chemical bonds.
Bond energy is the energy required to break one mole of a particular bond in the gaseous phase.
Formation of bonds releases energy (exothermic process), while breaking bonds requires energy (endothermic
process).
Intermolecular Forces
Definition
Intermolecular forces (IMFs) are the attractive or repulsive forces that occur between molecules. These
forces mediate interactions between individual molecules of a substance.
Intermolecular vs. Interatomic Forces
Feature Intermolecular Forces (IMFs) Interatomic Forces (Chemical Bonds)
Location Between separate molecules Within a molecule, between atoms
Strength Weaker (typically 0.4 to 40 kJ/mol) Stronger (typically 400 to 1000 kJ/mol)
Attractive forces due to sharing or transfer of
Nature Attractive or repulsive forces between molecular entities
electrons
Effect on Determine physical properties like boiling point, melting Determine molecular structure and chemical
State point properties
Hydrogen bonds, dipole-dipole interactions, London
Examples Covalent bonds, ionic bonds, metallic bonds
dispersion forces
Page 2
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Types of Intermolecular Forces
London Dispersion Forces (LDF):
Also known as van der Waals forces.
Present in all molecules, both polar and nonpolar.
Arise from temporary fluctuations in electron distribution, creating instantaneous dipoles.
Strength increases with the size and shape of the molecule.
Dipole-Dipole Interactions:
Occur between polar molecules that have permanent dipoles.
The positive end of one molecule is attracted to the negative end of another molecule.
Stronger than London dispersion forces.
Hydrogen Bonds:
A special type of dipole-dipole interaction.
Occur when a hydrogen atom bonded to a highly electronegative atom (such as nitrogen, oxygen, or
fluorine) is attracted to a lone pair of electrons on another electronegative atom.
Strongest type of intermolecular force.
Essential for many biological processes, such as DNA structure and enzyme function.
The Chemistry of Water
Water is a polar molecule due to the difference in electronegativity between oxygen and hydrogen atoms.
Water molecules form hydrogen bonds with each other, resulting in many unique properties, such as high
surface tension, high boiling point, and its ability to act as a universal solvent.
Water is essential for life and plays a crucial role in many biological and chemical processes.
Vectors in Two Dimensions
1.1 Introduction
In grade 10, you learned about vectors in one dimension. Now, we'll expand those concepts to two dimensions.
A vector has both a magnitude and a direction.
Vectors are often represented visually with arrows, where:
Length of the arrow indicates magnitude.
Arrowhead indicates direction.
When denoting a vector, an arrow is drawn over the symbol representing the physical quantity. Omitting the arrow
implies you're only referring to the magnitude.
Key Mathematics Concepts
Theorem of Pythagoras
Units and unit conversions
Equations
Trigonometry
Graphs
1.2 Resultant of Perpendicular Vectors
In grade 10, you learned about the resultant vector in one dimension, and now this will be extended to two
dimensions.
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The resultant vector represents the combined effect of multiple vectors acting simultaneously.
Vectors on the Cartesian Plane
In two dimensions, we use the Cartesian plane, which consists of two perpendicular axes: the x-axis (horizontal) and
the y-axis (vertical).
For example, a force F of magnitude 2 N acting in the positive x-direction can be drawn as a vector on the Cartesian
plane.
The length of the vector corresponds to the specified magnitude. The vector's starting point on the plane doesn't
affect the physical quantity, as long as the magnitude and direction remain the same. This is known as equality of
vectors.
Compass Directions
Compass directions can specify a vector's direction. The four cardinal directions are North, South, East, and West.
Directions can be combined if they are directly between two cardinal directions, such as North-East.
Bearing
A bearing is an angle, usually measured clockwise from North, used to specify direction.
This differs from the Cartesian plane, where angles are measured counter-clockwise from the positive x-direction.
The Resultant Vector
In Grade 10, you learned about adding vectors together in one dimension. The same principle can be applied for
vectors in two dimensions. Vectors that fall on the same line are called co-linear vectors.
To add co-linear vectors, use the tail-to-head method.
Worked example 1: Revision: head-to-tail addition in one dimension
Use the graphical head-to-tail method to determine the resultant force on a rugby player if two players on his team
are pushing him forwards with forces of → F1 = 600 N and → F2 = 900 N respectively and two players from the
opposing team are pushing him backwards with forces of → F3 = 1000 N and → F4 = 650 N respectively.
SOLUTION
Step 1: Choose a scale and a reference direction
Page 4
Atomic Combinations
Chemical Bonds
Chemical bonds are forces that hold atoms together to form molecules or compounds. These bonds arise from
the interactions between the positive nuclei and the negative electrons of atoms.
Atoms combine to achieve a more stable electron configuration, typically resembling that of a noble gas (8
valence electrons).
Types of Chemical Bonds
1. Covalent Bonds:
Formed by the sharing of electrons between two atoms.
Typically occur between two nonmetal atoms.
Examples: water (H₂O), methane (CH₄)
2. Ionic Bonds:
Formed by the transfer of electrons from one atom to another, resulting in the formation of ions.
Occur between a metal and a nonmetal.
The metal atom loses electrons to become a positively charged ion (cation), and the nonmetal atom gains
electrons to become a negatively charged ion (anion).
These oppositely charged ions are attracted to each other, forming the ionic bond.
Examples: sodium chloride (NaCl), magnesium oxide (MgO)
3. Metallic Bonds:
Found in metals.
Characterized by a "sea" of delocalized electrons that are free to move throughout the metallic structure.
This electron mobility accounts for the high electrical and thermal conductivity of metals.
Examples: copper (Cu), iron (Fe)
Molecular Shape
The shape of a molecule is determined by the arrangement of its atoms in three-dimensional space.
The shape influences a molecule's physical and chemical properties, such as its polarity, reactivity, and
biological activity.
Valence Shell Electron Pair Repulsion (VSEPR) theory helps predict the shape of molecules based on the
repulsion between electron pairs surrounding a central atom.
Basic Molecular Shapes:
Linear: Atoms are arranged in a straight line (e.g., CO₂)
Trigonal Planar: Three atoms are bonded to a central atom, forming a flat, triangular shape (e.g., BF₃)
Tetrahedral: Four atoms are bonded to a central atom, forming a three-dimensional pyramid shape (e.g., CH₄)
Bent: Similar to trigonal planar or tetrahedral but with one or more lone pairs of electrons, resulting in a bent
shape (e.g., H₂O)
Trigonal Pyramidal: Three atoms and one lone pair are bonded to a central atom, forming a pyramid shape (e.g.,
NH₃)
Page 1
, Created by Turbolearn AI
Electronegativity
Electronegativity is a measure of the ability of an atom in a chemical compound to attract shared
electrons in a covalent bond.
Electronegativity differences between atoms in a bond determine the polarity of the bond.
Linus Pauling scale is commonly used to represent electronegativity.
Electronegativity and Bond Polarity
Nonpolar Covalent Bond:
Occurs when two atoms have similar electronegativities.
Electrons are shared equally.
Example: H₂
Polar Covalent Bond:
Occurs when two atoms have different electronegativities.
Electrons are shared unequally, resulting in a partial positive charge (δ+) on the less electronegative atom
and a partial negative charge (δ−) on the more electronegative atom.
Example: H₂O
Ionic Bond:
Occurs when there is a large difference in electronegativity between two atoms.
Electrons are essentially transferred from one atom to another, resulting in the formation of ions.
Example: NaCl
Energy and Bonding
Energy is involved in the formation and breaking of chemical bonds.
Bond energy is the energy required to break one mole of a particular bond in the gaseous phase.
Formation of bonds releases energy (exothermic process), while breaking bonds requires energy (endothermic
process).
Intermolecular Forces
Definition
Intermolecular forces (IMFs) are the attractive or repulsive forces that occur between molecules. These
forces mediate interactions between individual molecules of a substance.
Intermolecular vs. Interatomic Forces
Feature Intermolecular Forces (IMFs) Interatomic Forces (Chemical Bonds)
Location Between separate molecules Within a molecule, between atoms
Strength Weaker (typically 0.4 to 40 kJ/mol) Stronger (typically 400 to 1000 kJ/mol)
Attractive forces due to sharing or transfer of
Nature Attractive or repulsive forces between molecular entities
electrons
Effect on Determine physical properties like boiling point, melting Determine molecular structure and chemical
State point properties
Hydrogen bonds, dipole-dipole interactions, London
Examples Covalent bonds, ionic bonds, metallic bonds
dispersion forces
Page 2
, Created by Turbolearn AI
Types of Intermolecular Forces
London Dispersion Forces (LDF):
Also known as van der Waals forces.
Present in all molecules, both polar and nonpolar.
Arise from temporary fluctuations in electron distribution, creating instantaneous dipoles.
Strength increases with the size and shape of the molecule.
Dipole-Dipole Interactions:
Occur between polar molecules that have permanent dipoles.
The positive end of one molecule is attracted to the negative end of another molecule.
Stronger than London dispersion forces.
Hydrogen Bonds:
A special type of dipole-dipole interaction.
Occur when a hydrogen atom bonded to a highly electronegative atom (such as nitrogen, oxygen, or
fluorine) is attracted to a lone pair of electrons on another electronegative atom.
Strongest type of intermolecular force.
Essential for many biological processes, such as DNA structure and enzyme function.
The Chemistry of Water
Water is a polar molecule due to the difference in electronegativity between oxygen and hydrogen atoms.
Water molecules form hydrogen bonds with each other, resulting in many unique properties, such as high
surface tension, high boiling point, and its ability to act as a universal solvent.
Water is essential for life and plays a crucial role in many biological and chemical processes.
Vectors in Two Dimensions
1.1 Introduction
In grade 10, you learned about vectors in one dimension. Now, we'll expand those concepts to two dimensions.
A vector has both a magnitude and a direction.
Vectors are often represented visually with arrows, where:
Length of the arrow indicates magnitude.
Arrowhead indicates direction.
When denoting a vector, an arrow is drawn over the symbol representing the physical quantity. Omitting the arrow
implies you're only referring to the magnitude.
Key Mathematics Concepts
Theorem of Pythagoras
Units and unit conversions
Equations
Trigonometry
Graphs
1.2 Resultant of Perpendicular Vectors
In grade 10, you learned about the resultant vector in one dimension, and now this will be extended to two
dimensions.
Page 3
, Created by Turbolearn AI
The resultant vector represents the combined effect of multiple vectors acting simultaneously.
Vectors on the Cartesian Plane
In two dimensions, we use the Cartesian plane, which consists of two perpendicular axes: the x-axis (horizontal) and
the y-axis (vertical).
For example, a force F of magnitude 2 N acting in the positive x-direction can be drawn as a vector on the Cartesian
plane.
The length of the vector corresponds to the specified magnitude. The vector's starting point on the plane doesn't
affect the physical quantity, as long as the magnitude and direction remain the same. This is known as equality of
vectors.
Compass Directions
Compass directions can specify a vector's direction. The four cardinal directions are North, South, East, and West.
Directions can be combined if they are directly between two cardinal directions, such as North-East.
Bearing
A bearing is an angle, usually measured clockwise from North, used to specify direction.
This differs from the Cartesian plane, where angles are measured counter-clockwise from the positive x-direction.
The Resultant Vector
In Grade 10, you learned about adding vectors together in one dimension. The same principle can be applied for
vectors in two dimensions. Vectors that fall on the same line are called co-linear vectors.
To add co-linear vectors, use the tail-to-head method.
Worked example 1: Revision: head-to-tail addition in one dimension
Use the graphical head-to-tail method to determine the resultant force on a rugby player if two players on his team
are pushing him forwards with forces of → F1 = 600 N and → F2 = 900 N respectively and two players from the
opposing team are pushing him backwards with forces of → F3 = 1000 N and → F4 = 650 N respectively.
SOLUTION
Step 1: Choose a scale and a reference direction
Page 4
