Scientists can classify matter as elements, compounds, or mixtures depending on the composition of an object (Ebbing & Gammon, 2011). When atoms with the same atomic number combine covalently or by metallic bonding, the resultant substance is an element. For instance, a molecule of oxygen contains two oxygen atoms that have undergone covalent bonding while copper metal contains copper atoms that are close together due to metallic bonding (Stoker, 2012). On the other hand, a compound refers to a chemical union of different elements, parts, or ingredients. Sodium chloride is a compound, which results from the chemical combination of sodium atoms and chlorine atoms. During the formation of compounds, elements lose their individuals properties with the release or absorption of significant amounts of heat energy. A mixture refers to a group of different materials combined by physical means (Stoker, 2012). The combination of sulfur powder and iron filings is an example of a mixture. Therefore, compounds and mixtures are different from one another based on the process of formation and the individual properties of the elements in the resultant substance. The formation of mixtures requires the release or absorption of heat energy while the formation of mixtures does not result in any energy change (Ebbing & Gammon, 2011). Elements in a mixture retain their physical and chemical properties. On the other hand, in compounds, elements undergo a change in their individual properties.
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A chemist can differentiate between a pure compound and a pure element by using various techniques. Investigation of the chemicals properties of a substance can reveal whether it is a compound or an element (Stoker, 2012). For instance, a chemist can investigate how the substance reacts with other substances, such as hydrogen, acids, oxygen, or alkalis. If a substance reacts with an acid with evolution of hydrogen gas, then the substance is metal, which is an element. Other elements, such as oxygen gas, react with metals other nonmetals to form metal oxides and nonmetal oxides respectively. Some compounds, such as metal carbonates, react with acids to form carbon dioxide, water, and salts. Chemists can also investigate the physical properties of the substances (Stoker, 2012). Some of the physical properties include melting and boiling point of substances. Elements have well defined boiling and melting points, which is not the case with compounds. X-ray studies can also enable chemists to tell whether a substance is an element or a compound. The x-ray spectrum identifies all elements that a substance contains (Ebbing & Gammon, 2011). For instance, a substance is an element if it contains only one element.
Compounds or molecules result from a chemical combination of atoms. Ionic bonds hold atoms in the structure of compounds, while covalent bonds hold atoms together in the structure of molecules. Compounds contain oppositely charged ions, whose attractions bound atoms together (Ebbing & Gammon, 2011). For, instance chlorine and sodium react to form sodium chloride, which is a compound with ionic bonding in its structure. Chlorine forms a negative ion while sodium forms a positive ion. Binding of atoms in molecules happens because the atoms share outermost electrons. Some atoms share electrons equally forming covalent bonds (Ebbing & Gammon, 2011). For instance, in the molecule of water, hydrogen atoms share electrons with oxygen atoms.
Ionic compounds result from the reaction between metals and nonmetals (Ebbing & Gammon, 2011). In the periodic table of elements, metals appear to the left hand side while the nonmetals appear to the right hand side. When metals and nonmetals react, metal atoms have a high likelihood of losing the outermost electrons while the nonmetals have a high likelihood of gaining electrons in the outermost shell. This is because the atoms of metals have fewer electrons in the outermost energy levels than the atoms of nonmetals do. Therefore, it is easier to lose the outermost electrons in metal atoms than losing the outermost electrons in the nonmetal atoms (Stoker, 2012). On the other hand, it is easier to acquire electrons in nonmetal atoms than is the case with metal atoms. Losing of the outermost electrons results in the formation of positive ions, while gaining of electrons in the outermost energy level results in the formation of negative ions (Moore, 2011). Therefore, the oppositely charged ions attract each other, which results in the formation of ionic bonds. An example of compounds that result from ionic bonding include sodium chloride, which consists of negatively charged chloride ions and positively charged sodium ions (Ebbing & Gammon, 2011). Another example is magnesium oxide, which consists of negatively charged oxygen ions and positively charged magnesium ions.
The nonmetals from the right hand side of the periodic table of elements possess relatively high electronegativities (Ebbing & Gammon, 2011). Therefore, the atoms of these elements have a higher ability to attract electrons than the metal atoms do. Examples of such elements include oxygen and sulfur, which react to form sulfur dioxide. Both the atoms of oxygen and sulfur have six electrons at the outermost energy level. However, oxygen atoms have smaller atomic radii than the atoms of sulfur do, and this explains why oxygen atoms have high electronegativities as compared to the atoms of sulfur (Stoker, 2012). Because oxygen atoms have higher electronegativities than sulfur atoms do, the shared pair of electrons tends to move closer toward the oxygen atoms and far from the sulfur atom, which leads to a partial positive charge on a sulfur atom and a partial negative charge on the oxygen atoms (Moore, 2011). The negative and positive charges attract one another, resulting in polar covalent bonding. Non-polar covalent bonding occurs when the atoms of the same nonmetal element join to from a molecule. For instance, two chlorine atoms may bond and form chlorine molecules. Sharing of the outermost electrons leads to stable octet structure, which is a characteristic of noble gases (Moore, 2011).