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Add information to the 4 sections corresponding to the the learning objectives. All must be completed.

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Bonding

• Understand ionic bonding:

o strong electrostatic attraction between oppositely charged ions

o effects ionic radius and ionic charge have on the strength of ionic bonding

o formation of ions in terms of electron loss or gain

o electronic configuration diagrams of cations and anions.

• Understand covalent bonding:

o strong electrostatic attraction between two nuclei and the shared pair(s) of electrons

between them

o dot and cross diagrams to show electrons in simple covalent molecules, including

those with multiple bonds and dative covalent (coordinate) bonds

o the relationship between bond lengths and bond strengths in covalent bonds

o tetrahedral basis of organic chemistry.

• Understand metallic bonding:

o de-localised electrons

o positive metal ions

o regular layer structure.

• Understand the following intermolecular forces:

o van der Waals

o dipole-dipole

o hydrogen bonding.

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a metal and a non metal

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Metallic bonding is a type of chemical bonding that rises from the electrostatic attractive force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions.

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Metallic bonding is a type of chemical bonding that rises from the electrostatic attractive force between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions.

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only metallic atoms

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an example would be a chloride ion, the atom chlorine gains an electron to become the ion chloride

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an example would be a chloride ion, the atom chlorine gains an electron to become the ion chloride

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Ionic bonding is the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion.

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Ionic bonding is the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion.

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covalent compounds have no free electrons and no ions so they don't conduct electricity.

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Intermolecular attractions are attractions between one molecule and a neighbouring molecule. The forces of attraction which hold an individual molecule together (for example, the covalent bonds) are known as intramolecular attractions.

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force of attraction

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the non metals share electrons to complete a full outer shell

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the non metals share electrons to complete a full outer shell

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For example, the covalent bond, involving sharing electron pairs between atoms, is much stronger than the forces present between neighbouring molecules.

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For example, the covalent bond, involving sharing electron pairs between atoms, is much stronger than the forces present between neighbouring molecules.

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When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms, forming ions. The resulting compound is called an ionic compound.

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When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms, forming ions. The resulting compound is called an ionic compound.

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electrons are transferred from one atom to another to form charged atoms, which are called ions

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In each of these reactions, the metal atoms give electrons to the non-metal atoms. The metal atoms become positive ions and the non-metal atoms become negative ions.

There is a strong electrostatic force of attraction between these oppositely charged ions, called an ionic bond.

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positive ions are formed from atoms losing electrons and negative ions are formed by atoms gaining electrons. The reason why atoms become charged after adding or losing electrons is because they now have an imbalance in the number of protons and electrons. These two sub-atomic particles have a negative charge (electron) and a positive charge (proton), so when atoms have an equilibrium in the number of protons and electrons, there is no charge.

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Oxygen gains two electrons to have an octet. The ionic bond between ions results from the electrostatic attraction of opposite charges. The final formula of iron oxide is Fe ₂ O ₃.

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a common example of an ionic bond that forms when chlorine and sodium transfer electrons leaving sodium with a positive charge and chlorine with a negative one

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the reason why atoms like to lose or gain electrons is because they all want a full outer shell, just like the noble gases. For the metals, they tend to lose electrons to gain a positive charge. It is easier to lose electrons when there are about less than 4 electrons in the outer shell since that would require less energy, compared to gaining electrons from another atom as that would take up a lot more energy.

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A metallic bond is a type of chemical bond formed between positively charged atoms in which the free electrons are shared among a lattice of cations. In contrast, covalent and ionic bonds form between two discrete atoms. Metallic bonding is the main type of chemical bond that forms between metal atoms.

Metallic bonds are seen in pure metals and alloys and some metalloids. For example, graphene (an allotrope of carbon) exhibits two-dimensional metallic bonding. Metals, even pure ones, can form other types of chemical bonds between their atoms. For example, the mercurous ion (Hg ₂ ²⁺) can form metal-metal covalent bonds. Pure gallium forms covalent bonds between pairs of atoms that are linked by metallic bonds to surrounding pairs.

The outer energy levels of metal atoms (the s and p orbitals) overlap. At least one of the valence electrons participating in a metallic bond is not shared with a neighbor atom, nor is it lost to form an ion. Instead, the electrons form what may be termed an "electron sea" in which valence electrons are free to move from one atom to another.

The electron sea model is an oversimplification of metallic bonding. Calculations based on electronic band structure or density functions are more accurate. Metallic bonding may be seen as a consequence of a material having many more delocalized energy states than it has delocalized electrons (electron deficiency), so localized unpaired electrons may become delocalized and mobile. The electrons can change energy states and move throughout a lattice in any direction.

Bonding can also take the form of metallic cluster formation, in which delocalized electrons flow around localized cores. Bond formation depends heavily on conditions. For example, hydrogen is a metal under high pressure. As pressure is reduced, bonding changes from metallic to nonpolar covalent.

Because electrons are delocalized around positively-charged nuclei, metallic bonding explains many properties of metals.

Electrical Conductivity - Most metals are excellent electrical conductors because the electrons in the electron sea are free to move and carry charge. Conductive nonmetals (e.g., graphite), molten ionic compounds, and aqueous ionic compounds conduct electricity for the same reason - electrons are free to move around.

Thermal Conductivity - Metals conduct heat because the free electrons are able to transfer energy away from the heat source and also because vibrations of atoms (phonons) move through a solid metal as a wave.

Ductility - Metals tend to be ductile or able to be drawn into thin wires because local bonds between atoms can be easily broken and also reformed. Single atoms or entire sheets of them can slide past each other and reform bonds.

Malleability - Metals are often malleable or capable of being molded or pounded into a shape, again because bonds between atoms readily break and reform. The binding force between metals is nondirectional, so drawing or shaping a metal is less likely to fracture it. Electrons in a crystal may be replaced by others. Further, because the electrons are free to move away from each other, working a metal doesn't force together like-charge ions, which could fracture a crystal through the strong repulsion.

Metallic Luster - Metals tend to be shiny or display metallic luster. They are opaque once a certain minimum thickness is achieved. The electron sea reflects photons off the smooth surface. There is an upper frequency limit to the light that can be reflected.

The strong attraction between atoms in metallic bonds makes metals strong and gives them high density, high melting point, high boiling point, and low volatility. There are exceptions. For example, mercury is a liquid under ordinary conditions and has a high vapor pressure. In fact, all of the metals in the zinc group (Zn, Cd, Hg) are relatively volatile.

Because the strength of a bond depends on its participant atoms, it's difficult to rank types of chemical bonds. Covalent, ionic, and metallic bonds may all be strong chemical bonds. Even in molten metal, bonding can be strong. Gallium, for example, is nonvolatile and has a high boiling point even though it has a low melting point. If the conditions are right, metallic bonding doesn't even require a lattice. It has been observed in glasses, which have an amorphous structure.

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Learn about metallic bonding, including the definition, its properties, and examples of materials that exhibit it.

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covalent bonding is also another way to try and gain a full outer shell. the electrons are shared, meaning that the atoms are connected to each other for this to happen. This type of bonding an happen between the same atoms or between different elements or compounds.

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All of these substances have very strong covalent bonds between the atoms, but much weaker forces holding the molecules together. When one of these substances melts or boils, it is these weak 'intermolecular forces' that break, not the strong covalent bonds.

Because the weak intermolecular forces break down easily these substances have low melting and boiling points. This means simple molecular substances are gases, liquids or solids with low melting points, and low boiling points.

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these are large structures where the atoms are all linked to each other, and the way that they are connected is with strong covalent bonds. That, therefore, means that giant covalent structures are extremely hard and have high melting and boiling points.

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The reason why these structures have high melting and boiling points is due to the requirement of a high amount of energy to break these bonds. The stronger the structure, the higher the amount of energy is needed to change the state of an element or compound.

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Metals conduct electricity because the electrons are free to move and carry charge.

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Metals conduct electricity because the electrons are free to move and carry charge.

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Intermolecular forces are the forces which mediate interaction between molecules, including forces of attraction or repulsion which act between molecules and other types of neighbouring particles, e.g., atoms or ions.

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In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole or multipole moment.

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BBC - Higher Bitesize Chemistry - Patterns in the periodic table : Revision, Page4

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A BBC Bitesize secondary school revision resource for Higher Chemistry on patterns in the periodic table: boiling points, atomic size, ionisation energy.