Unit 8: Bonding & IMF's |
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Demonstrations
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covalent bonding-hydrogen
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covalent bonding-Fluorine
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ionic bonding-NaF
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IMF's
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IMF's
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Videos
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Regents Core Curriculum
- The outermost electrons in an atom are called the valence electrons. In general, the number of valence electrons affects the chemical properties of an element. (3.1l)
The placement or location of elements on the Periodic Table gives an indication of physical and chemical properties of that element. The elements on the Periodic Table are arranged in order of increasing atomic number. (3.1y) - Elements can be classified by their properties and located on the Periodic Table as metals, nonmetals, metalloids (B, Si, Ge, As, Sb, Te), and noble gases. (3.1v)
- Elements can be differentiated by their physical properties. Physical properties of substances, such as density, conductivity, malleability, solubility, and hardness, differ among elements. (3.1w)
- Elements can be differentiated by chemical properties. Chemical properties describe how an element behaves during a chemical reaction. (3.1x)
- Some elements exist in two or more forms in the same phase. These forms differ in their molecular or crystal structure, and hence in their properties. (5.2f)
- For Groups 1, 2, and 13-18 on the Periodic Table, elements within the same group have the same number of valence electrons (helium is an exception) and therefore similar chemical properties. (3.1z)
The succession of elements within the same group demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, metallic/nonmetallic properties. (3.1aa) - The succession of elements across the same period demonstrates characteristic trends: differences in atomic radius, ionic radius, electronegativity, first ionization energy, metallic/nonmetallic properties. (3.1bb)
- A chemical compound can be represented by a specific chemical formula and assigned a name based on the IUPAC system. (3.1cc)
- A solution is a homogeneous mixture of a solute dissolved in a solvent. The solubility of a solute in a given amount of solvent is dependent on the temperature, the pressure, and the chemical natures of the solute and solvent. (3.1oo)
- Two major categories of compounds are ionic and molecular (covalent) compounds. (5.2g)
- Chemical bonds are formed when valence electrons are: transferred from one atom to another (ionic), shared between atoms (covalent), mobile within a metal (metallic) (5.2a)
- In a multiple covalent bond, more than one pair of electrons are shared between two atoms. (5.2e)
Molecular polarity can be determined by the shape of the molecule and the distribution of charge. Symmetrical (nonpolar) molecules include CO2, CH4, and diatomic elements. Asymmetrical (polar) molecules include HCl, NH3, and H2O. (5.2l) - When an atom gains one or more electrons, it becomes a negative ion and its radius increases. When an atom loses one or more electrons, it becomes a positive ion and its radius decreases. (5.2c)
- When a bond is broken, energy is absorbed. When a bond is formed, energy is released. (5.2i)
- Atoms attain a stable valence electron configuration by bonding with other atoms. Noble gases have stable valence configurations and tend not to bond. (5.2b)
Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces. These properties include conductivity, malleability, solubility, hardness, melting point, and boiling point. (5.2n) - Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions. (5.2d)
- Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. Electronegativity values are assigned according to arbitrary scales. (5.2j)
- The electronegativity difference between two bonded atoms is used to assess the degree of polarity in the bond. (5.2k)
Intermolecular forces created by the unequal distribution of charge result in varying degrees of attraction between molecules. Hydrogen bonding is an example of a strong intermolecular force. (5.2m)
