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Atoms and Chemical Bonding

Science, Grade 6

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Table Of Contents: Atoms and Chemical Bonding

1. Models of the Atom

2.1. Atomic Models
Scientists use models to explain things that cannot be seen directly. Since atoms are too small to be seen, scientists use models to describe their structure. The first model of the atom was proposed by the Greek philosopher, Democritus, who hypothesized that all matter was made of small particles, called atoms. He thought that different materials were made of atoms of different shapes and sizes. Over time, as new information has been discovered through experimentation, the model of the atom has changed to incorporate the new observations.
2.2. Dalton's Atomic Theory
John Dalton was an educator who was very interested in science. In the early 1800s, Dalton conducted experiments on gases. From his results he formed a model known as atomic theory. Dalton hypothesized the following: 1) All matter is made of atoms and atoms are small particles that cannot be created, divided or destroyed. 2) All atoms of the same element are identical and different elements have different types of atoms. 3) Atoms combine with other atoms to form new substances. Much of Dalton's atomic theory is still accepted today.
2.3. Thomson's Model of the Atom
In 1897, J.J. Thomson performed experiments with currents of electricity inside empty glass tubes, which disproved Dalton's theory that atoms cannot be divided. He discovered that atoms have negatively charged subatomic particles, which he called "corpuscles." These particles are now called electrons.Thomson assumed that atoms were neutral, so if an atom had negatively charged particles, they must have an equal number of positively charged particles. In 1904, Thomson proposed a model of the atom as a solid sphere with equal numbers of positive and negative charges spread throughout, much as raisins might be embedded in the surface of a pudding.
2.4. Rutherford's Experiment
In 1911, Ernest Rutherford and his students, conducted his "gold-foil" experiment to confirm Thomson's model of the atom. He fired positively charged alpha particles from a radioactive source at a piece of gold foil. Rutherford thought that if Thomson's model was correct, then the mass of the atom was spread out throughout the atom and there would be very little to deflect the alpha particles. As Rutherford expected, most alpha particles traveled right through the gold foil with little deviation, but to his surprise a few alpha particles deflected almost directly backwards. The results of this experiment led to the creation of a new model of the atom. This new atomic model consisted of a positively charged nucleus, containing most of the atomic mass of the atom, and negatively charged electrons orbiting around the nucleus like planets around the Sun.
2.5. Bohr's Model of the Atom
The Rutherford atomic model had a major drawback; it could not explain how the electrons stayed in a stable path around the nucleus. In Rutherford's model of the atom, the electrons are scattered around the nucleus. However, the positively charged nucleus would attract the negatively charged electrons. The force of attraction would pull the electrons towards the nucleus, making the atom collapse. In 1913, Niels Bohr made changes to Rutherford's model to explain the stability of atoms. Bohr proposed electrons move in stable, or stationary, orbits at fixed distances from the nucleus. Each orbit had an energy associated with it. The closest orbit had the lowest energy and the energy increased with the distance from the nucleus. If an electron moves between orbits, then energy in the form of light (or photons), is absorbed or emitted.
2.6. The Modern Atomic Theory
Many scientific advances have continued to change the model of the atom. In 1926, Erwin Schrˆdinger, an Austrian physicist, took the Bohr atomic model one step further. Schrˆdinger used mathematical equations to describe the likelihood of finding an electron in a certain position. This atomic model is known as the Quantum Mechanical model of the atom. Unlike the Bohr model, which defines the exact path of an electron, the Quantum Mechanical model predicts the probability of the electron's location. This model can be represented as a nucleus surrounded by an electron cloud. The probability of finding the electron is greatest where the cloud is most dense and least likely in a less dense area of the cloud. Until 1932, the atom was believed to be composed of a positively charged nucleus surrounded by negatively charged electrons. In 1932, James Chadwick discovered the neutron, an uncharged particle contained in the nucleus. Since then, through continued experimentation, more particles have been discovered in the atom. Quarks are believed to be even smaller units than protons and neutrons. In turn, quarks are made up of vibrating strings of energy. The search to find even smaller particles that make up an atom continues.

2. Pause and Interact

3.1. Review
Use the whiteboard text tool to answer the questions.
3.2. Scientists and their model of the atom
Drag the name of the scientist to their model of the atom.

3. Atomic Configuration and Bonding

4.1. Electron Configuration
The specific way the electrons are arranged in an atom is called the electron configuration. Electrons play an important role in how elements interact with each other and form compounds. Electrons are negatively charged particles that surround the nucleus in the form of a cloud. The atomic number of an atom, which is equivalent to its number of protons, also represents the number of electrons in that atom. The electrons are distributed among orbital shells or energy levels (1, 2, 3 and so on) that are different distances from the nucleus. The larger the number of the energy level, the farther it is from the nucleus. Electrons that are in the highest or outmost energy level are called valence electrons. The valence electrons are the ones that are lost, gained or shared during chemical bonding. In the electron configuration for oxygen, which has a total of 8 electrons, the first orbital or energy level closest to the nucleus is filled with two electrons. The second energy level can hold up to eight electrons. It begins to fill with the remaining six electrons when the first level is full.
4.2. Valence Electrons and Bonding
The electrons in the outermost shell are called the valence electrons. The outermost orbital shell, called the valence shell, is most often involved in chemical bonding. Elements in the same group in the periodic table have the same number of electrons in their valence shell. For example, all elements in group 1, alkali metals, have 1 valence electron. Group 1 atoms prefer to lose 1 electron to become stable. However, all elements in group 17, halogens, have 7 valence electrons. These atoms will gain 1 electron to fill their valence shell. Atoms with 4 or more electrons in the valence shell like to gain electrons to fill the shell. Atoms with 3 or less electrons in their valence shell like to lose electrons to reveal the full inner shell. Atoms with full valence shells will not combine with other elements.
4.3. Electron Dot Diagram
The number of valence electrons in an atom of an element determines many properties of that element, including the ways in which the atom can bond with other atoms. An electron dot diagram is often used to depict the valence electrons in an atom. Each atom of an element has a specific number of valence electrons, ranging from 1 to 8. The electron dot diagram includes the element symbol, surrounded by dots. Each dot represents 1 valence electron. The dots are spaced out above, below, to the left, and to the right of the symbol for the first 4 valence electrons. For atoms with greater than 4 valence electrons, the dots must be paired up. The dot diagrams for atoms can also be used to show the bond between different atoms in a molecule.
4.4. Stability of Atoms
Atoms of most elements are more stable and are less likely to react with other atoms, when they have 8 valence electrons in their outer shell. For example, atoms of neon, argon, krypton, and xenon are very unreactive because they all have 8 valence electrons. Atoms usually react in a way that makes each atom more stable by losing, gaining, or sharing electrons in a chemical bond with other atoms.
4.5. Chemical Bonds
In 1916, Gilbert Newton Lewis, an American scientist, proposed that chemical bonds are formed due to the electron interaction between atoms. His work established the basis of what we know today about chemical bonding. Atoms combine with other atoms through chemical bonds, which result from the strong attractive forces that exist between the atoms. Atoms bond together to become more stable by having a full valence shell. Certain elements are more reactive than others and will be more likely to bond with other elements. There are three main types of chemical bonding: covalent bonding, ionic bonding and metallic bonding. The bond between 2 nonmetals is usually a covalent bond. Whereas the bond between a metal and nonmetal atom is an ionic bond. Metal atoms bond by forming a metallic bond.

4. Pause and Interact

5.1. Review
Use the whiteboard text tool to fill in the table.
5.2. Electron Configuration
Select the best answer(s) and then click on the check button.

5. Ionic Bonding

6.1. Charged Particles
A normal atom has a neutral charge with an equal number of protons and electrons. For example, a neutral atom of sodium has eleven protons and eleven electrons. However, since atoms are more stable with a full valence shell, an atom can gain or lose electrons to become stable. If a sodium atom loses 1 electron, it will have a full valence shell but it will no longer be neutral. The atom will now have 11 protons and 10 electrons for a net charge of +1. The neutral atom becomes charged when there are an unequal number of protons and electrons.
6.2. Forming Ions
A neutral atom that becomes charged by gaining or losing electrons is called an ion. An atom can acquire a positive or negative charge depending on whether the number of electrons is greater or less than the number of protons in the atom. An atom with more electrons than protons, has a negative charge and is a negative ion, or anion. However, an atom with more protons than electrons, has a net positive charge and is a positive ion, or a cation. Since an electron and a proton have equal but opposite unit charges, the charge of an ion is always expressed as a whole number of unit charges and is either positive or negative. For example, calcium loses 2 valence electrons, to form a positive ion with a +2 net charge which is expressed as Ca²⁺. Alternatively, an oxygen atom gains two electrons to form a negative ion with a -2 charge, which is expressed as O²⁻. Molecules can also be ions, known as polyatomic ions. Most polyatomic ions are anions with one notable exception, the ammonium cation (NH₄⁺).
6.3. Ionic Bonds
Ionic bonds form as a result of the attraction between positive and negative ions. Ionic bonding normally occurs between metal atoms and nonmetal atoms. Atoms with partially filled outer shells are unstable. To become stable, the metal atom loses one or more electrons in its outer shell, forming a positively charged ion or cation with a stable electron configuration. These electrons are then gained by the nonmetal atom, causing it to form a negatively charged ion or anion, also with a stable electron configuration. The attraction between the oppositely charged ions causes them to come together and form an ionic bond. In this example, sodium (Na) and chlorine (Cl) ions are attracted to each other in a 1:1 ratio and combine to form sodium chloride (NaCl), common table salt.
6.4. Ionic Compounds
Atoms that form ionic bonds are called ionic compounds. Ionic compounds are formed from a positive cation interacting with a negative anion. The negative and positive ions form a repeating pattern. This pattern creates a crystal lattice which is orderly and three dimensional. The lattice is generally hard and brittle due to the strong bonds between the ions. The repeating pattern also provides the substance with a high melting point because of the high energy required to break all the atomic bonds. When an ionic compound is added to water, the water molecules pull apart the ions of the solid lattice. The resulting solution contains charged ions that conduct electricity.
6.5. Naming Ionic Compounds
The names of ionic compounds are written by listing the name of the positive ion followed by the name of the negative ion. Therefore, a series of rules is needed to name the positive and negative ions before we can name these compounds. Single atomic positive ions have the name of the element from which they are formed. However, since some metals form positive ions in more than one oxidation state, (Ex. Fe²⁺, Fe³⁺ ) the charge on the ion is indicated by a Roman numeral in parentheses immediately after the name of the element (Ex. iron(II), iron(III)). Polyatomic (more than two atoms) positive ions often have common names ending with the suffix "-onium" such as hydronium (H₃O⁺) or ammonium (NH₄⁺). Negative ions that consist of a single atom are named by adding the suffix "-ide" to the stem of the name of the element. The name of polyatomic negative ions usually ends in either "-ite" or "-ate." The "-ite" ending indicates a low oxidation state (nitrite ion NO₂⁻), whereas the "-ate" ending indicates a high oxidation state (nitrate ion NO₃⁻). Oxidation state shows the total number of electrons that an atom has gained or lost in order to form a chemical bond with another atom.
6.6. Writing Formulas for Ionic Compounds
When an ionic compound is formed, the ions must combine in a way that the total charges of the compound equal zero. For example, let's write the correct formula for magnesium chloride. The first step is to write the formulas for the cation Mg²⁺ and anion Cl⁻. Next, drop the positive and negative signs; crisscross the superscripts so that they become subscripts and reduce when possible by finding the least common multiple. In this example, the two chlorine ions, with a total charge of -2, balance the +2 charge of the magnesium ion. The cation is always listed first before the anion, resulting in the formula MgCl₂.

6. Pause and Interact

7.1. Review
Use the whiteboard tools to name the following compounds.

7. Covalent and Metallic Bonding

8.1. Covalent Bonds
Covalent bonding is a type of chemical bonding that occurs between two non-metallic atoms. It is characterized by the sharing of one or more pairs of electrons between atoms. By sharing electrons, two atoms can mutually complete their valence shells to become more stable. For example, since each chlorine atom has 7 electrons in its outer shell, two chlorine atoms will each share an electron to obtain a complete outer shell and form a stable Cl₂ molecule. For every pair of electrons shared between two atoms, a single covalent bond is formed. Atoms can also share two or three pairs of electrons and are named accordingly as double and triple bonds. A single line indicates a bond between two atoms, double lines indicate a double bond and triple lines represent a triple bond.
8.2. Nonpolar and Polar Covalent Bonding
There are two types of covalent bonding - nonpolar and polar. Nonpolar bonding results when two identical non-metals equally share electrons between them. Diatomic molecules such as O₂ or I₂ form nonpolar covalent bonds where both atoms share the electrons equally. Polar bonding results when two different non-metals unequally share electrons between them. Compounds such as carbon dioxide, ammonia, and water have polar covalent bonds. Certain other compounds, such as ethane (C₂H₆), have both polar and nonpolar bonds. Ethane, has polar bonds between the carbon and hydrogen, and nonpolar bonds between the two carbon atoms.
8.3. Covalent Compounds
The atoms in covalent compounds, also known as molecular compounds, are bonded together by covalent bonds. Unlike ionic compounds which form a regular pattern, covalent compounds form individual molecules that are not connected to each other. Due to weak intermolecular forces, most covalent molecules or covalent compounds are liquids or gases at room temperature, with low melting and boiling points. And since covalent molecules do not separate into ions when dissolved in water, they are poor conductors of electricity. Although most covalent compounds are gases or liquids at room temperature, there's a class of solid compounds known as covalent network solids that are bonded by covalent bonds, but in a lattice structure. Such compounds are typically hard, transparent, and have high melting points. Examples include diamond, quartz and graphite, among others.
8.4. Covalent Formulas and Names of Covalent Compounds
The formula for a covalent compound can be derived from its name by writing the symbols for the first and second element and translating the prefixes into subscripts. For example, sulfur trioxide would be written as SO₃. Covalent compounds are named in a similar manner to binary ionic compounds. (Binary compounds are compounds made up of only two elements). To name binary covalent compounds apply the following rules: 1) Name the first nonmetal using the element's full name. 2) Name the second nonmetal element with an "-ide" ending as if it were an anion. 3) Indicate the number of atoms present (its subscript) of each element with the prefixes, "mono-", "di-", "tri-" and so on. For example, the formula name for CCl₄, with one carbon, and four chlorine atoms is carbon tetrachloride.
8.5. Metallic Bonding
The force that holds atoms together in a metallic substance is a metallic bond. A metallic substance consists of closely packed atoms, arranged in a very compact and orderly pattern. The valence shells (outermost electron shell) of the metal atoms overlap with many of the neighboring atoms. As a result, instead of orbiting their atoms, the valence electrons leave individual atoms and continually move throughout the metal structure from one atom to another. The atoms that the electrons leave behind become positive ions, and the interaction between such ions and valence electrons provides the binding force that holds a metallic structure together. This free movement of the valence electrons also provides metals a number of their unique characteristics, such as strength, malleability, ductility, thermal and electrical conductivity, opacity and luster.

8. Pause and Interact

9.1. Review
Use the whiteboard tools to name the following compounds.
9.2. Chemical Bonding
Select the best answer(s) and then click on the check button.

9. Vocabulary Review

10.1. Atoms and Chemical Bonding Vocabulary Matching Review

10. Virtual Investigation

11.1. Atoms and Chemical Bonding
How does the structure of the atom determine its ability to bond with other atoms? In this virtual lab you will build a model of an atom for several elements. You will be given the atomic number and atomic mass of the element. Recall that the number of protons in an atom is equal to the atomic number of the element. The number of neutrons is the atomic mass minus the atomic number. The number of electrons in a neutral atom is equal to the number of protons. You will use the Bohr model of the atom to place the electrons in the correct energy level, or orbital. The first energy level holds a maximum of two electrons; the second holds a maximum of eight electrons. From the model of the atom, you will determine whether or not it is possible for the atoms of the element to bond with other atoms.

11. Assessment

12.1. Atoms and Chemical Bonding
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